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& J. Lambert 

A New 


ourth Edition 

Group number 








6 939 

9 012 







22 990 

24 312 







39 102 

40 080 





54 938 
































88 905 






102 905 





























137 340 

138 910 






192 200 


















































10 811 
























28 086 

30 974 

32 064 


39 948 



















58 710 

63 540 

65 370 


72 590 

































107 870 

112 400 


118 690 




131 300 






























200 590 



208 980 





















Rod on 











Atomic weight 



Atomic number 



140 907 

144 240 











59 rnium 





232 038 


238 030 















[ ] This is the mass number of the Isotope with the longest Known half life of the element Indicated. 
* Thit Is the most number of the most sloble or best known Isotope of the element indicated 


157 250 


162 500 

164 930 


168 934 


174 970 











































101 ium 



A New 


School Certificate Chemistry 

Graded Problems in Chemistry to Ordinary Level 

(with ami without answers) 

Worked Examples and Problems in Ordinary Level Chemistry 

A Class Book of Problems in Chemistry to Advanced Level 

(with and without answers) 

A Simple Approach to Atomic Theory 

Problems and Worked Examples in Chemistry to Advanced Level 

The Essentials of Qualitative Analysis 


The Essentials of Volumetric Analysis 


Intermediate Organic Chemistry 

Inorganic and Physical Chemistry 

A Simple Approach to Organic Chemistry 

Revision Notes in Advanced Level Chemistry 

Volume One. Organic Chemistry 

Volume Two. Inorganic Chemistry 

Volume Three. Physical Chemistry 


Tables for Elementary Analysis 


Practical Chemistry 

A New 


A. Holderness m.Sc. f.r.i.c 

Formerly Senior Chemistry Master at 
Archbishop Holgate's Grammar School, York 


John Lambert m.Sc. 

Formerly Senior Chemistry Master at 
King Edward's School, Birmingham 



Hcinemann Educational Books Ltd 





ISBN 435 64408 4 (cased edn) 
ISBN 435 64410 6 Gimp edn) 

©A. Holderness and J. Lambert 1961, 1964, 1969 
First Published 1961 
Second Edition 1961 

Reprinted 1962 

Third Edition 1964 

Reprinted 1965, 1966, 1968 

Fourth Edition 1969 

Reprinted 1970 

Published by 

Heinemann Educational Books Ltd 

48 Charles Street, London W1X 8AH 

Printed in Great Britain by 

Butler and Tanner Ltd, Frome, Somerset 

Preface to the Fourth Edition (1969) 

The developments now taking place constitute the greatest upheaval 
in school chemistry teaching in living memory. The problems it pre- 
sents arc as great in text-book authorship as in teaching, including as 
they do the implications of changing nomenclature and units as well 
as great modifications of syllabuses. 

In the matter of nomenclature, we have adopted the Stock nota- 
tion for metallic compounds as well as the systematic naming of such 
groups as the oxides of nitrogen. At the same time, it has to be 
remembered that 'trivial' usages such as ferrous, ferric and cupric 
are found in all the English chemical literature before the 1960's and 
should be familiar to users of it. Consequently, these now obsolete 
forms are occasionally mentioned along with the revised usages. 
Similarly, everyday 'trivial' terms such as red lead, lime, caustic soda, 
bicarbonate and others are mentioned in their place. 

The SI system of units has been adopted, e.g., cm 8 and dm 3 for 
volume measurements, except in a short section devoted to normality 
(for historical reasons) when the litre (1) and ml units are retained. 
Energy change is stated in terms of AH and kJ (with the equivalent 
calorie figure at times). The mole unit is used frequently, though 
with some trepidation about the ambiguity that could arise if our 
young chemists use the term loosely. Terms such as gram-molecule 
and gram-atom (though declared obsolete in some quarters) are 
retained since they appear in the syllabus requirements of some 
Boards. The banning of the term equivalent will probably come as a 
jolt to older teachers, especially when it is substituted by an ex- 
pression such as a mole of±H t SOi. Enough of the historical develop- 
ment of atomic theory has been retained to enable pupils to under- 
stand that ia C = 12 is only a minor variation of convenience on the 
older H = 1. The important experimental work formerly devoted to 
equivalents has been adapted to determine formulae of simple com- 
pounds and verify suitable equations. Atomic weight determination 
is left to the mass spectrometer (with treatment suited to O-level). 

The syllabuses of the examining Boards have, of course, a large 
area of common ground but show some variations on the fringes. 
The general emphasis of all Boards on such topics as periodicity, 



electronic usages, molecular shapes, metallic bonding, the Avogadro 
constant and its relation to the mole, the pH scale, elementary 
chromatography, macromolecules (carbon, metals, starch, nylon), 
kinetic theory and elementary organic chemistry, has been met by 
the introduction or necessary expansion of these topics. In addition, 
the fringe variations, e.g., London's requirements of polystyrene 
and perspex, some lithium and coppcr(l) chemistry and calorimetnc 
measurements, or the Joint Matriculation Board's of the co-ordinate 
bond and some social and economic considerations in chemistry, 
have been taken into account. The new requirements of the Hong 
Kong Certificate of Education have also been satisfied by the in- 
clusion of its special topics, e.g., electroplating, oxidation number, 
lead accumulator, Leclanche cell and the behaviour of hydrogen 
chloride in toluene. It is hoped that, by this means, all the substantial 
requirements of all the Boards have been met, including multiple 
choice questions of which a practice set is supplied. 

November 1969 

A. H. 
J. L. 

Preface to the First Edition 

A New Certificate Chemistry is designed as an alternative to the same 
authors' School Certificate Chemistry. The twenty-three years which 
have passed since the appearance of School Certificate Chemistry 
have brought inevitable changes in the science of Chemistry and in 
its presentation at middle-school level. Many of these changes have 
been covered by alterations in the text of the book from time to 
time, but the electronic outlook which has developed in recent years 
has made more decisive change and a new title desirable. 

All the topics and many of the characteristics of the older book 
have been retained since they have proved, over the years, very 
acceptable to teachers. Among these characteristics are the insistence 
on adequate detail in experimental work, a full presentation of 
elementary molecular theory and chemical arithmetic, and plenty of 
material for examination revision. 

The principal features which justify the title, A New Certificate 
Chemistry, are the treatment of oxidation-reduction and the relation 
between metals and non-metals in electronic terms, wherever suit- 
able. Similarly, topics such as the acid-base relation and 
neutralisation are presented in ionic terms. These changes have 
involved the introduction of ionic equations in much greater numbers 
than before. At the same time, it has not been forgotten that ionic 
equations do not lend themselves to calculations involving actual 
compounds encountered in practical work. To meet this situation, 
molecular equations have been retained in sufficient numbers and in 
suitable contexts, so that firm contact is retained between equations 
and calculations and the actual chemical materials which appear in 
the stock bottles, the beakers and the filter-papers used by pupils in 
chemistry. To this extent, A New Certificate Chemistry is a mixture 
of the old and the new in middle-school chemistry. This, we hope, is 
not unsuitable to the present transitional stage of development in 
school science. 

The authors have pleasure in acknowledging the very helpful 
comments they have received, while preparing this book, from some 
practising science teachers, more particularly Messrs A. B. Adamson, 




F. E. P. Alford, F. W. Ambler, H. G. Andrew, H. C. Cockroft, 
E. H. Coulson, J. B. Guy, I. G. Jones, P. N. Lawrence, E. W. Moore, 
T. A. Muir and J. Turpie. 

They would also like to thank Mr E. Dickinson for his assistance 
and advice at proof stage, and the eight School Certificate Examina- 
tion Boards for permission to use their questions. 

University of Bristol (B.). University of Cambridge, Local (C). 

Central Welsh Board (C.W.B.). University of Durham (D.). 

University of London (L.). Joint Matriculation Board (N.U.J.B.). 

Oxford Local (O.)Oxford and Cambridge (O. and C). 

March 1961 

A. H. 
J. L. 


Preface to the fourth edition 


Preface to the first edition 


Chemical nomenclature 


Physical and chemical change; elements, compounds, 



Atomic theory 


Laws of combination by weight (experimental) 


Atomic weight 


The behaviour of gases under temperature and 




Experimental work on formula, reactions 

and equivalents 


Calculations involving weights 


Atomic structure; shape of molecules; 





Molecular theory 


Formulae of gases 


Volumetric analysis 


Electrochemical series and electrolysis 


Types of reactions 


Air, combustion and rusting 


Water and solution 


Acids, bases and salts 










Hydrogen peroxide and ozone 


Carbon and its oxides; flame 


Carbonates and hydrogen carbonates 


Organic chemistry 


Silicon and its compounds 




Hydrogen chloride and the chlorides 


Bromine and iodine 




Hydrogen sulphide and the sulphides 


Oxides and acids of sulphur; sulphates 


Nitrogen and ammonia 




33 Nitric acid and the nitrates 435 

34 Oxides of nitrogen 445 

35 Phosphorus 455 

36 Metals; extraction and uses; non-metals compared 463 

37 Radioactivity; a brief account 498 

38 Revision Notes 509 
Definitions 524 
General questions (including multiple choice) 528 
Answers 531 
Approximate atomic weights 532 
Logarithmic tables 533 
International atomic weights 535 
Index 537 

List of Plates 

1 The Concorde 

2 Atmosphere above a town 

3 (a) Water pipe blocked by deposits 
(b) Stalagmites and stalactites 

4 Atomic hydrogen used in welding 

5 (a) Manufacture of carburetted water-gas 
(b) Ammonia convertors 

6 Manufacture and use of fertiliser 

7 Blown glass 

8 (a) Oslo evaporator 

(b) Open-hearth furnace 

facing page 18 












Chemical Nomenclature 

THE word 'nomenclature' means 'system or scheme of naming'. 
The following simple treatment of chemical nomenclature will 
enable you to state the composition of most chemical substances 
directly from their names. 


The commonest elements have been known for a long time and 
there is generally no system about their naming. Recently isolated 
melals, however, have been given names ending in -ium or -urn, e.g., 
radium, platinum, osmium, duminium, while recently named non- 
metals have been given names ending in -on, e.g., argon, xenon. 


Binary compounds 

The name-ending -ide is given to compounds containing only two 
elements 1 and the nature of the elements is indicated in the two words 
of the name, e.g., copper oxide, CuO; hydrogen sulphide, H 2 S. 

The number of atoms of one of the elements contained in a mole- 
cule of the compound is sometimes indicated by a prefix to the second 
part of the name, e.g., carbon (//sulphide, CS 2 ; phosphorus rr/chloride, 
PC1 3 ; carbon te/rachloride, CC1 4 ; phosphorus pentoxide, P 2 6 . 

Acids and salts 

A great many acids contain hydrogen, oxygen and a third element, 
e.g., H,S0 4 , HN0 3 , H 3 P0 4 . The commonest and most stable of 
such acids is usually highly oxidised and to it is given a name which 
ends in -ic and is derived from the element it contains in addition to 
hydrogen and oxygen, e.g., sulphuric acid, H 2 S0 4 , and nitric acid, 
HN0 3 . An acid containing the same elements but less oxygen has the 
name-ending changed to -ous, while one with less oxygen still takes 

1 Hydroxides are exceptions to this rule but in these compounds the three 
elements present are indicated in the name. Another exception is an acid 
salt of hydrogen sulphide, e.g. sodium hydrogen sulphide, NaHS, but here 
again the name is self-explanatory. In salts like ammonium chloride, NH.CI, 
the NH, group has been treated as if it were an element. 




the prefix hypo- with the -ous ending. An acid with a higher proportion 
of oxygen than the -ic acid takes the prefix per- with the -ic ending. 

The corresponding salts have names of the form -ate, -ite, hypo 
- -ite and per — ate. Thus : 

perchloric acid, HC10 4 . 
chlor/c acid, HCIO3. 
chlorous acid, HC10 2 . 
hypochlorous acid, HCIO. 

potassium perchlorate, KC10 4 . 
potassium chlorate, KC10 3 . 
potassium chlorite, KC10 2 . 
potassium hypochlorite, KCIO. 

Notice also the following common pairs of acids and salts: 


sulphuric acid, H 2 S0 4 . 
sulphurous acid, H 2 S0 3 . 
nitric acid, HN0 3 . 
nitrous acid, HNO a . 

sulphates, e.g., Na 2 S0 4 . 
sulphites, e.g., Na 2 S0 3 . 
nitrates, e.g., KN0 3 . 
nitrites, e.g., KN0 2 . 

A useful rule to remember is that a salt with the name ending -ate 
or -ite usually contains three elements, one of which is oxygen, e.g., 
lead sulphate, PbS0 4 , copper nitrate, Cu(N0 3 ) a , sodium hypo- 
chlorite, NaClO. Note the two compounds, ferrous sulphide, FeS 
(two elements), and ferrous sulphate, FeS0 4 (three elements includ- 
ing oxygen). These two are frequently confused by beginners. 

Recent recommendations of the Chemical Society and the Inter- 
national Union of Pure and Applied Chemistry have made certain 
chemical names 'trivial' and they are to go out of use. Examples are 
names in -ic and -ous derived from names of metals, e.g., ferric and 
ferrous, cupric and cuprous. The recommended replacement takes the 
form of the ordinary name of the metal with its operative valency 
stated in brackets in Roman numerals. Where necessary, the majority 
of the names in the present edition have been changed to conform to 
this notation. Some cases of the old usages have been left because 
they occur in chemical literature before about 1960 and students 
should be familiar with them. Examples of change are the following: 

Cuprous oxide 
Cupric oxide 
Ferrous sulphate 
Ferric sulphate 
Plumbous chloride 
Plumbic chloride 
Manganese dioxide 

CopperfJ) oxide, Cu 2 0. 
Copper(II) oxide, CuO. 
Iron (II) sulphate, FeS0 4 . 
Iron(IlI) sulphate, Fe^SO^. 
Lead(Il) chloride, PbCl 2 . 
Lead(IV) chloride, PbCl 4 . 
Manganese(IV) oxide, Mn0 2 . 

Chapter 1 

Physical and Chemical Change 
Elements, Compounds, Mixtures 


THE science of Chemistry sets before itself, as its primary objects, 
first, the determination of the nature and properties of the non- 
living matter which surrounds us in that portion of the crust of the 
earth to which we have access, and secondly, the preparation of new 
substances, scientifically interesting or generally useful, from the 
materials which Nature has provided. In trying to determine the 
nature of substances, chemists have been greatly interested in the 
changes which these undergo when subject to conditions which they 
normally do not encounter— high temperature, high pressure, 
extreme cold, contact with other materials under varying conditions, 
and so on. It is largely from the changes which materials undergo 
when subject to these conditions that chemists, particularly in the last 
150 years or so, have drawn conclusions about their nature. 

But changes are multifarious. Some of them do not interest the 
scientist, as such, at all. The change wrought when the hand of an 
artist transfers paint from palette to canvas and, spreading it in a 
particular way, creates a work of art, is a matter of jesthetics, not of 
science. Any one change may be viewed from many different angles. 
When iron rusts, a chemist is concerned with the different properties 
which the iron and rust possess. How does each react with acids, 
alkalis and other reagents? He also tries to give an explanation of what 
occurred to the iron when it rusted. The physicist will want to know 
whether the density, the conductivity and the specific heat of the rust 
are the same as that of iron, from which it has been made. An 
economist thinks of the huge cost which accompanies the change, for 
millions of pounds are spent yearly in an endeavour to prevent iron 
from rusting. It may be that what the chemist and the physicist find 
out about the rusting of iron will help the economist (or, more 
directly, the manufacturer of iron articles). It is mainly as a result of 


chemical research that rustless and stainless steels have appeared, 
while during that research, the physicist has been careful to ensure 
that the elasticity and tensile strength of the steel have not been im- 
paired by the process which made them rustless. Clearly, we must 
attempt to define, if only roughly, the kind of changes in which the 
chemist is interested and to which the name 'chemical' can properly 
be applied. With the object of attaining some kind of definition of 
'chemical changes', we will now examine a few changes in the hope 
that, from them, some conclusions may emerge. 

1. Hold one end of a piece of 
magnesium ribbon in tongs and 
put the other end in a Bunsen 
flame. Note the intense brilliance 
of the flame of the burning mag- 
nesium and the nature of the 
residue — a white ash — which re- 

2. Take a small piece of sodium 
in tongs from the oil under which 
it is kept, and, never touching it 
with your fingers, cut it into pieces 
about the size of a very small pea. 
Drop these pieces in turn on to the 
surface of a little distilled water in a 
small beaker. Note how the sodium 
melts into a ball, darts about the 
surface of the water, produces a 
hissing sound and finally disappears 
with a small flash and explosion. 
Heat the resulting clear liquid on 
an iron dish until no more steam 
is given off. On cooling, a white 
solid is left. If added to water, this 
solid dissolves but does not show 
the same vigorous action as sodium. 
It is a new substance, caustic soda. 

3. Heat some roll sulphur on a 
deflagrating spoon. Note how the 
sulphur melts and later begins to 
burn with a blue flame. It gradually 
decreases in amount and finally the 
spoon will be left empty. The 
sulphur has not simply been anni- 
hilated. Its disappearance is due to 
its conversion into a new gaseous 
substance which is invisible, but 
whose presence in the air can be 
detected merely by its irritating 

Repeat the experiment with 
platinum by holding a loose coil of 
platinum wire in a Bunsen flame. 
Note the white-hot glow of the 
metal, but contrast its unchanged 
appearance, after cooling, with the 
white ash left by the magnesium. 

To distilled water in a beaker, 
add some common salt and stir the 
mixture. The common salt under- 
goes an obvious change; it gradu- 
ally disappears forming a solution 
and being no longer visible as a 
white solid. Put the liquid into a 
porcelain dish and heat gently 
until all the water has evaporated 
off. The common salt reappears in 
its original white solid form. 

Powder some roll sulphur in a 
mortar, then heat it gently in a test- 
tube, shaking all the time. Notice 
how the sulphur melts to an amber- 
coloured liquid (other changes will 
occur if it is more strongly heated) 
and that this liquid, on cooling, 
returns to its original condition as a 
yellow solid. 


smell or, if you prefer some less 
commonplace evidence, by burning 
the sulphur in a gas-jar and adding 
to the jar some blue litmus solution. 
The gas, sulphur dioxide, will turn 
it red. 

A little consideration will quickly show that the six changes we 
have considered above are not all the same in nature. They fall, 
actually, into two classes, which are distinguished by the following 

1. All the changes in the right-hand column were easily reversible; 
the molten sulphur returned to the solid form when cooled; the 
platinum wire ceased to glow and regained its original appearance 
when removed from the flame; the common salt was recovered by 
evaporating off the water. Contrast these results with those of the 
changes in the left-hand column. The white ash from magnesium 
was totally unlike the original magnesium, and it would be a difficult 
matter to obtain magnesium from it; on evaporation of water in 
Experiment 2, we recovered not sodium, but caustic soda, from which 
sodium cannot easily be obtained; the sulphur became part of a gas 
from which it would be difficult to recover sulphur. These changes 
are not reversible. The easily reversible type of change is called 
'physical change*; the more permanent type 'chemical change'. 

2. In none of the physical changes recorded in the right-hand 
column was a new kind of matter formed ; we began with platinum, 
common salt and sulphur and, after the change, finished with just 
those materials. In all the chemical changes recorded in the left-hand 
column, some new kind of matter was formed: magnesium was con- 
verted to the white powdery ash, magnesium oxide, sodium to caustic 
soda and solid sulphur to the gas, sulphur dioxide — a new kind of 
matter each time. This is characteristic of chemical change. 

3. The physical changes in the right-hand column were not accom- 
panied by any marked external effects. The solution of common salt 
in water and the melting of sulphur were not violent changes. The 
action of sodium with water, however, produced enough heat to melt 
the sodium and was violent enough to be slightly explosive at the 
end; the burning of magnesium produced intense heat and light, and 
the burning of sulphur similar, but less intense, effects. The chemical 
changes were the more violent, and were accompanied by heat 
changes. This is commonly the case. 

4. We carried out no weighings during our experiments, but it 
could actually have been shown that, in all three physical changes 
which we recorded on the right, no change of weight occurred; the 
sulphur, platinum and common salt weighed just as much before the 


changes as after them. In the three chemical changes, however, it 
would have been found that the white ash weighed more than the 
magnesium, the caustic soda more than the sodium and the gaseous 
sulphur dioxide more than the sulphur. (These gains in weight are 
made at the expense of other materials which lose in weight corre- 
spondingly; the gains in the case of magnesium and sulphur were 
made at the expense of the air and, in the case of sodium, at the 
expense of the water.) 

We thus distinguish two kinds of changes — chemical changes and 
physical changes. 

Now consider the following suggestions about a few common 
changes and decide by comparison with those discussed above 
whether the changes are physical or chemical. The correct classi- 
fication appears on p. 5. 

1. (a) Melting of ice; (b) conversion of water to steam. 

Are the changes easily reversed? Are there any noticeably violent 
external effects? 

2. Burning of coal. 

Does the coal appear to weigh the same as the products after burn- 
ing it? (Appearances here are deceptive, see p. 25.) Can we easily 
obtain coal again from its products of combustion? Are there any 
noticeable external effects while coal is burning? 

3. Rusting of iron. 

Can iron be easily recovered from the rust? 

4. Magnetising iron. 

Can the iron be readily de-magnetised? Are there any marked 
changes during magnetisation? 

5. A coal-gas explosion. 

Is this change violent? Is there considerable heat change? 

6. Heating of the filament of an electric light globe by the current. 
Is the filament readily cooled again? Does it appear changed when 

cooled ? 

7. The melting of candle-wax. 

Is the liquid wax easily solidified again? Docs it then appear the 
same as the original wax? Are there any marked heating effects as 
the melting occurs? 

We may now summarise the characteristics of chemical and 
physical change in the table, p. 5. 



1. Is generally easily reversible 

2. Produces no new kind of 

3. Is not accompanied by great 
heat change (except latent 
heat effects accompanying 
changes of state) 

4. Produces no change of weight 


1. All cases of the melting of a 
solid to a liquid (or the re- 

2. All cases of vaporisation of a 
liquid (or the reverse) 

3. Magnetisation of iron 

4. The heating of a metal wire 
by electricity 


1. Is generally not easily re- 

2. Always produces a new kind 
of matter 

3. Is usually accompanied by 
considerable heat change 

4. Produces individual sub- 
stances whose weights are diff- 
erent from those of the original 
individual substances 

Thus, if two substances, A 
and B, react chemically and 
are changed into substances 
C and D, the weight of C 
will be different from the 
weight of A or B, and the 
weight of D will be different 
from the weight of A or B. 


1. The burning of any substance 
in air 

2. The rusting of iron 

3. The slaking of lime 

4. Explosion of coal-gas or hy- 
drogen with air 


Elements and Compounds 

We will begin our study of elements and compounds by consider- 
ing two fairly simple chemical changes. 

1. Put a little red oxide of mercury (mercuric oxide) into a dry test- 
tube. Heat it, rotating the test-tube so that it does not become mis- 
shapen. A silvery mirror gradually appears on the upper part of the 
test-tube (where it is cool), and later silvery globules of mercury will 
be seen. When the mirror begins to appear, insert a glowing splint 
of wood into the test-tube. It is rekindled. This is because the in- 
visible gas, oxygen, is coming off from the heated oxide. 


It is clear that, under the action of heat, mercuric oxide has yielded 
two products —mercury and oxygen. 

2. Repeat the above experiment, using lead nitrate. Brown fumes 
are given off in this case (they are called nitrogen dioxide), and, by 
the test given above, it can be shown that oxygen is also liberated. 
Finally a yellow solid will remain in the test-tube. This solid is 

These experiments show that both mercuric oxide and lead nitrate 
must be fairly complex substances. This is obvious from the fact that 
mercuric oxide yielded, under the action of heat, two substances, 
mercury and oxygen, while lead nitrate yielded three — litharge, 
nitrogen dioxide and oxygen. The question now arises whether 
these products can themselves be split up further into still simpler 
substances. The answer to this question is that in two cases they 
can; from litharge we can, by suitable chemical means, obtain lead 
and oxygen, and from nitrogen dioxide, nitrogen and oxygen. This 
means that litharge and nitrogen dioxide are themselves complex 
substances. How much further can this process of splitting up into 
simpler products be carried? Can we obtain from the lead, nitrogen, 
oxygen and mercury, into which we have resolved our original lead 
nitrate and mercuric oxide, any substances which are simpler still? 
The answer now is that we cannot. By no chemical process whatever 
is it possible to obtain from lead, mercury, oxygen or nitrogen any 
substance simpler than themselves. Clearly, these four simple sub- 
stances are different from the more complex mercuric oxide and lead 
nitrate. The number of substances which like lead, oxygen, mercury 
and nitrogen, are incapable of being split up into simpler substances 
is small. There is very good reason to believe that about a hundred 
such substances may exist, and, of them, nearly all are actually known 
on the earth. To them the name 'elements' has been given. We may 
now define this term. 

Definition. An element is a substance which cannot by any known 
chemical process be split up into two or more simpler substances. 

A list of the elements is given on p. 535. We may here mention a 
few of the commoner ones. All the metals — lead, zinc, iron, copper, 
tin, platinum, gold, silver and the rest — are elements; so also are the 
oxygen and nitrogen of the air, together with carbon, sulphur, 
phosphorus, iodine and others to the number of over one hundred. 
Remember the characteristic they all possess - they cannot by any 
known chemical process be made to yield substances simpler than 

From this small band of elements, all other substances on the earth 
are made. The number of chemical substances known is more than a 
million and a half. All of these, except the elements themselves, are 


made up of two or more elements combined together. They are 
called 'compounds'. It is astounding to reflect that all compounds 
on the earth, from the simplest, which, like water, contain only two 
elements, to those complex materials of which our own bodily tissues 
are composed, are made from about one hundred simple, elementary 
materials. The elements are indeed a small select band. 

There follows a short list of common compounds and the elements 
which compose them. 



Common salt 
Blue vitriol 
Oil of vitriol 

Oxygen; hydrogen 
Oxygen; hydrogen; carbon 
Sodium; chlorine 
Potassium; nitrogen; oxygen 
Calcium; oxygen; carbon 
Copper; sulphur; oxygen; hydrogen 
Hydrogen; sulphur; oxygen 
Silicon; oxygen 
Aluminium; silicon; oxygen; hydrogen 

(Note how commonly oxygen occurs; it is the most widely dis- 
tributed of all the elements.) 

We may now define a compound. 
Definition. A compound is a substance which contains two or more 
elements chemically combined together. 

We have found it necessary to use the expression 'chemically com- 
bined'. The meaning of it is connected with the idea of chemical 
change which was discussed earlier. We must now try to obtain a 
clearer idea of the meaning of the expression and, to do this, we shall 
contrast the properties of mixtures and compounds in the work of 
the next section. 

Mixtures and Compounds 

Weigh out 28 g of iron filings and 16 g of sulphur (any multiples 
or fractions of these weights will do equally well). Grind the two 
thoroughly in a mortar and put about half of the mixture into a dry 
test-tube. Heat the test-tube, at the bottom, with a small flame. The 
mixture will glow. When it docs so, remove the flame, and hold the 
test-tube over a mortar as a precaution against breakage. The glow 
will then spread slowly through the mixture without further heating. 
Allow the test-tube to cool, then break it away from the mass of 
material left. A dark grey, almost black, solid will be found. 

In the following experiments, we shall compare the properties of 


the original mixture of iron and sulphur with those of the black solid 
left after heating it. 


I. Action with 

2. Action of a 


Action with 



4. Action of 
dilute hydro- 
chloric acid 

Mixture before heating 

Place enough of the 
mixture in a test-tube to 
fill about 2-5 cm of its 
depth. Half fill the test- 
tube with water, shake it 
well, then allow the test- 
tube to stand. The denser 
iron will settle more 
rapidly than the lighter 
sulphur and form a layer 
below it. The experiment 
separates the iron from 
the sulphur. 

Rub one end of a bar 
magnet well into the mix- 
ture, raise it and tap 
gently. The iron filings 
will have been attracted 
by the magnet and will 
adhere to it. The sulphur 
will not. They are sepa- 

Half fill a test-tube with 
carbon disulphide, add 
some of the mixture and 
shake for a few minutes. 
Filter the mixture through 
a dry filter-paper and 
funnel on to a dry clock- 
glass. Allow the filtrate to 
evaporate to dryness (in a 
fume-chamber) at the or- 
dinary temperature. On 
the clock-glass, a yellow 
deposit of solid sulphur 
will be left. The carbon 
disulphide has dissolved 
out the sulphur from the 
mixture and so separated 
it from the iron. 

Add dilutchydrochloric 
acid to some of the mix- 
ture in a test-tube. Warm 
gently. There is rapid 
effervescence. Apply a 
lighted taper to the test- 

Solid left after heating 
Carry out the same test. 
The solid settles as a 
single layer with no sign 
of separation of the iron 
from the sulphur. 

Repeat with the other 
end of the magnet. A very 
little iron (left unattacked 
by the sulphur) may be 
attracted by the magnet 
but it will be very much 
less than before. The bulk 
of the iron is not attracted 
from the black solid and 
is not separated from the 

Repeat the experiment 
described opposite. A 
very slight deposit of 
sulphur (left unattacked 
by the iron) may remain, 
but the great bulk of the 
sulphur has not been 
separated from the iron. 

Repeat the experiment 
described opposite. Effer- 
vescence occurs again. 
Apply the following two 
tests to the gas: 

1. Smell very cau- 




Mixture before heating 
tube. The resulting slight 
explosion shows that the 
gas is hydrogen. The iron 
has reacted with the acid 
to produce this gas. The 
sulphur remains un- 

Solid left after heating 

tiously. The rather dis- 
gusting smell is similar to 
that of rotten eggs. 

2. Apply a lighted 
taper to the test-tube. The 
gas burns with a blue 
flame but without explo- 
sion. It iS HYDROGEN 

Reason for the above Differences of Properties 

It is clear that the solid left after heating the mixture of iron and 
sulphur differs greatly in properties from the original mixture itself. 

1. Before the heating, the iron could be separated from the sulphur 
by physical methods. For example, by shaking with water, we 
took advantage of the physical property of density to separate the 
denser iron from the lighter sulphur; we also separated the iron from 
the sulphur by using its physical magnetic properties and by physical 
solution of the sulphur in carbon disulphide. In none of these experi- 
ments was any chemical action involved, but these physical methods 
could not separate the sulphur from the iron after the mixture had 
been heated. Physical methods of separation were then useless. 

2. Again, in the mixture before heating, the two elements clearly 
exercised their own independent properties. The iron was attracted 
by the magnet just as it would have been if the sulphur had not been 
present; the sulphur dissolved in carbon disulphide without inter- 
ference from the iron. Similarly, during the action of dilute hydro- 
chloric acid on the mixture, the iron reacted with the acid exactly as 
if no sulphur were present, while the sulphur itself remained un- 
changed. After the heating, however, the black solid left showed 
properties of its own. The sulphur present in it was no longer dis- 
solved out by carbon disulphide nor was the iron attracted by a 
magnet, while the separate densities of the two elements were no 
longer available for use in their separation. The action of dilute 
hydrochloric acid gave an entirely different reaction, with evolution 
of hydrogen sulphide instead of hydrogen. So we see that during the 
heating the separate properties of the iron and sulphur were lost and 
the new properties of the black solid, ferrous sulphide, appeared. 

The reason for this difference is that before the heating the two 
elements were simply mixed together, while during the heating they 
underwent chemical combination, forming the compound, ferrous 
sulphide. As a result of this change, the elements were united by a 
chemical link or bond instead of being merely close together in space. 



The nature of this bond has been the object of much speculation, and 
it is now known to be electrical. 

3. Another characteristic difference between physical mixing and 
chemical combination is apparent from this experiment. During the 
mixing of the iron and sulphur in the mortar, no change was observed 
except a kind of averaging of the colours of the two elements so that 
the mixture had a colour between the grey of iron and the yellow of 
sulphur. During the chemical combination, however, enough heat 
was given out to raise the whole mass to a bright red glow, once the 
action had been started by the external application of heat. Chemical 
combination is often accompanied by heat changes of this kind, but 
physical mixing is not. 

4. A further and most important difference between mixtures and 
chemical compounds is that the composition of a compound by 
weight is fixed and unalterable, while that of a mixture may vary 
within wide limits. For example, pure ferrous sulphide always con- 
tains the iron and sulphur in the proportion of 56 g of iron to 32 g 
of sulphur, and no variation from this proportion is ever found. 
Mixtures of iron and sulphur may, however, have any desired com- 

These differences between compounds and mixtures are summar- 
ised below. 


(a) The constituents can be sepa- 
rated from one another by 
physical methods. 

(b) Mixtures may vary widely in 

(c) Mixing is not usually accom- 
panied by external effects such 
as explosion, evolution of heat 
or volume change (for gases). 

((/) The properties of a mixture 
are the sum of the properties 
of the constituents of the mix- 


The constituent elements cannot 
be separated by physical methods; 
chemical reactions are necessary. 
Compounds are absolutely fixed 
in their compositions by weight. 
Chemical combination is usually 
accompanied by one or more of 
these effects. 

The properties of a compound are 
peculiar to itself and are usually 
quite different from those of its 
constituent elements. 


1. What are considered to be the main distinctions between a chemical 
compound and a mixture? Explain why the liquid obtained by mixing 
sodium chloride with water is not regarded as a chemical compound. (O. 

2. When a piece of sodium is placed in water it diminishes in size 
gradually, and finally disappears. In what way is the disappearance of the 
sodium different from the ordinary process of solution in water? Give 



experiments in support of your views. How could metallic sodium be 
recovered from the liquid? (N.U.J.B.) 

3. What is the essential difference between a chemical and a physical 

Indicate clearly the chemical and physical changes involved in the follow- 
ing processes, giving full reasons in each case: (a) the addition of metallic 
sodium to water; (/>) the solution of sodium chloride in water; (c) the 
heating of magnesium in air; (</) the heating of ammonium chloride; (e) the 
addition of water to concentrated sulphuric acid. (L.) 

4. Describe the experiments you would carry out in seeking to determine 
whether a given white powder is a pure substance or a mixture. If the 
substance is a pure chemical compound, how would you propose to 
ascertain whether it is (i) a salt; (ii) a basic oxide, or (iii) a peroxide? 

5. Illustrate three differences between metallic and non-metallic 
elements by reference to the properties of iron and sulphur. Describe 
three tests by means of which you would prove that the compound, iron 
sulphide, formed by heating a mixture of iron filings and sulphur, differs 
from the original mixture. (N.U.J.B.) 

Chapter 2 

Atomic Theory 

WE have seen in the last chapter that there are about 100 kinds 
of simple matter called elements, and that all other kinds of 
matter have been formed by the chemical combination of two or more 
of these elements. Thinkers have speculated for centuries in an en- 
deavour to elucidate further the structure of matter, and, more 
particularly, to shed light on the structure of the elements. The ancient 
Greek thinker, Democritus (about 400 B.C.), began the speculations, 
the Roman Lucretius (about 350 years later) took up the question, 
and, following these two ancient philosophers, there has appeared 
a succession of thinkers, European and Arabian, whose speculations 
culminated at the beginning of the nineteenth century in the ideas of 
an Englishman, John Dalton, of Manchester. His suggestions won 
universal acceptance for themselves and lasting scientific fame for 
their author. 


Dalton's life-time abounded in famous names and exciting hap- 
penings — the French Revolution, Austerlitz, Trafalgar, Nelson, 
Napoleon, Wellington, Waterloo — but the thoughts of this Quaker, 
slowly maturing as he pursued an obscure and uneventful existence, 
have proved more potent in their influence on human modes of living 
in the succeeding century than all the wars and alarums of his day. 
His Atomic Theory is the foundation of modern chemistry. 

Dalton's love of precision and truth is illustrated by the following 
story concerning him. Dalton had given a course of lectures, and at 
the end a student came to him with the request for a certificate of 
attendance. The great chemist looked up his records and found that 
the student had missed one lecture during the course. Dalton refused 
to sign the attendance certificate, but, after considering a few minutes, 
he said, *If thou wilt come tomorrow I will go over the lecture thou 
hast missed.' Having quietened his conscience in this manner over 
the missed lecture, Dalton presumably signed the certificate. 




Ideas about atoms 

You will have seen, from the work of the last chapter, that, if 
ordinary powdered sulphur is mixed with water, it does not dissolve 
and can be removed by filtration. That is, the sulphur particles are too 
large to pass through the pores of the filter-paper. If, however, we 
dissolve sulphur in carbon disulphide and filter the liquid, filtration 
does not separate the sulphur. That is, the sulphur particles are now 
small enough to pass through the pores of the filter-paper. 

If sulphur is very finely ground and mixed with water, filtration 
does not remove all the sulphur. Some is fine enough to remain 
permanently mixed with the water in what is called a colloidal solution. 
When this liquid is examined by the ultra-microscope, it can be seen 
that sulphur particles are in rapid random motion, which continues 
indefinitely at room temperature. This motion is now known to be 
caused by collisions between the minute sulphur particles and par- 
ticles (molecules) of water. The movement is shown by all colloidal 
particles in water. It was first observed (1827) by the botanist, Brown, 
for particles of pollen in water and is known as the Brownian 

This idea that elements are capable of being split up into very tiny 
particles was first evolved by Greek and Roman thinkers from 
another angle. They said that if we were to take a piece of, say, gold 
and cut it up into small pieces, and cut those pieces into smaller 
pieces, and those pieces into small pieces, and so on, a time would 
ultimately come when the dividing process would have to stop. The 
tiny particles of gold which we had then obtained would be in- 
capable of being divided any further; they would be the smallest 
possible particles of gold which could ever be obtained. The Greeks 
gave them the name 'atoms'. As a temporary, though incomplete, 
definition of an atom we may say that it is the smallest, indivisible 
particle of an element. If the Greeks had been Englishmen they 
might have called these particles 'indivisibles', for the word 'atom' 
meant to a Greek what 'indivisible' means to an Englishman. We 
still call them 'atoms' today. 

Hypothesis and theory 

The idea that elements are made up of atoms is called the atomic 
theory. This word 'theory' does not often mean very much to a 
beginner in Chemistry, so let us be quite clear about it. A scientific 
theory is a scientific idea which was thought of by somebody, 
suggested by him in a scientific book or journal, and accepted by 
other scientists after due consideration. 'So-and-so's theory' means 
'So-and-so's accepted idea'. The process of getting an idea accepted 



may be a long one; there will be arguments, objections, improve- 
ments of the idea, but, if it finally wins acceptance by scientists 
generally, it will be called a theory. When the idea is first put forward, 
and is still in the 'argument-and-objection' stage, it is called a hypo- 
thesis; later, if generally accepted, a theory. 


Dalton's Atomic Theory was first put forward in 1808. It soon 
gained general acceptance and then stood, virtually unchanged, for 
about a century. Discoveries made early in the twentieth century, 
however, showed that the theory must be modified. For this reason, 
the Atomic Theory will first be considered in its original, nineteenth- 
century form; the changes made in it will be mentioned later, in 
Chapter 9. 

The Atomic Theory, we have seen, goes back to the Greeks, yet we 
always speak today about Dalton's Atomic Theory. There is good 
reason for this. The reason is that, while the Greeks put forward the 
idea that atoms exist, they did no more. They left the idea vague and 
untested. Dalton changed this vague imagining into a set of concrete 
suggestions about atoms which could be tested by experiment. This 
change from vagueness to precision and experimental test justifies 
his claim to the theory. Below are given the ideas which together 
make up the Atomic Theory of Dalton (1808). 

The Atomic Theory states: 

1. Matter is made up of small, indivisible particles called atoms. 

2. Atoms are indestructible and they cannot be created. 

3. The atoms of a particular element are all exactly alike in every 
way and are different from the atoms of all other elements. 

4. Chemical combination takes place between small whole num- 
bers of atoms. 

These ideas are so important that we shall discuss all except the 
first more fully. 

Atoms are indestructible and cannot be created 

The important aspect of this idea is that, by chemical action, it is 
possible to alter only the state of combination of a number of atoms, 
not to reduce their number or add to it. If we start a chemical 
reaction with, say, a thousand million atoms of hydrogen, then we 
shall finish that reaction with exactly a thousand million, neither 
more nor less. They may have altered their state of combination— 
they may have become, for example, free hydrogen gas instead of 

dalton's atomic theory 


being combined with oxygen in water— but the same number will be 

The atoms of a particular element are all exactly alike in every way 
and are different from the atoms of all other elements 

The most important point is that this statement includes in it the 
idea that all atoms of the same element are exactly alike in mass, but 
are different in mass from the atoms of any other element. The 
theory said that if we collected together, say, one thousand atoms of 
sulphur from all corners of the earth, every one of those atoms of 
sulphur would be exactly the same as every other. The same would 
be true of any number of atoms of copper. But the mass of each 
sulphur atom would be different from the mass of each copper atom. 

Be sure you understand the universality of this idea. Consider, say, 
oxygen. Oxygen occurs in hundreds of thousands of compounds — 
water, sugar, litharge, blue vitriol, alcohol, oil of vitriol, starch, and 
so on. If we were to collect one oxygen atom from each of these 
hundreds of thousands of compounds, every one of those oxygen 
atoms would, said the Atomic Theory, be absolutely and completely 

It is clear that we could not test the theory in this way because the 
atoms are so small that we could not examine them or weigh them 
even if we could separate them. 

Idea on size of atoms 

It is difficult to have any conception of the size of atoms. The 
following diagram may help you to understand how small an atom 
really is. 

One gram of hydrogen liberated here and 
allowed to diffuse uniformly over the sur- 
face of the earth. 

This column of air (1 
cm* CTOSS-section), any- 
where on the surface of 
the earth, would contain 
over 100,000 of those 
hydrogen atoms. (It is 
assumed they are allowed 
to diffuse freely.) 

Fig. I. 
Idea on size of atoms. 

In other words, there are over 100,000 times more atoms in 1 g of hydrogen, 
than square centimetres on the whole of the surface of the earth. 



Chemical combination takes place between small whole numbers of 

It follows from the supposition that atoms are indivisible that they 
must combine in whole numbers. Dalton made the additional as- 
sertion that these whole numbers are small. By this he meant that 
atoms commonly combine in such numbers as 3 atoms of one 
element with 1 atom of another, or 2 atoms of one element with 5 
atoms of another, or 1 atom of one element with 1 atom of another. 
Cases such as 67 atoms of one element combining with 125 atoms of 
another, or 322 atoms of one element with 27 atoms of another, were 
unknown. The numbers of atoms combining together are almost 
always small, though, of course, in any one laboratory experiment 
there will be millions of exactly similar combinations taking place. 

We must now emphasise the fact that what has been 
stated above is a set of ideas. In science, ideas are treated 
with scant respect unless they can be backed up by experi- 
mental results. None of the statements made above can be 
tested by direct observation. We cannot line up a thousand 
oxygen atoms and inspect them to see if they are all alike, 
or count the number of hydrogen atoms which combine with 
one atom of oxygen to see if the number is small. We must 
deduce, from the Atomic Theory, some results which ought 
to follow from it, and which can be put to experimental 

To find an experimental test of the Atomic Theory's assertion that 
atoms are indestructible and cannot be created 

If both these statements are true, it follows that the total number of 
atoms present at the end of a chemical reaction must be the same as 
the number present at the beginning of it, though they will be 
differently combined. That is, the total mass of the products of a 
chemical action should be the same as the total mass of the starting 
materials. Experiment shows this to be true and the situation is 
expressed in the Law of Conservation of Matter. Experiments suit- 
able for illustrating it are given on p. 24, where the Law is formally 

To find an experimental test of the Atomic Theory's assertion that all 
atoms of the same element are exactly alike but different from atoms 
of other elements 

It is obvious from their different properties that atoms of different 
elements differ from one another. 
If atoms of the same element are all identical, it follows that all 



pure samples of the same chemical compound must be identical in 
composition by mass, however they are prepared. For example, 
samples of black copper oxide, containing one atom of copper and 
one atom of oxygen in each molecule, should have the same com- 
position by mass, irrespective of the sources of copper and oxygen. 
This can be tested by preparing copper oxide in several different ways 
and analysing the various samples. Accurate experiment shows this 
identity of composition in all the samples. The facts are stated in the 
Law of Constant Composition (or Definite Proportions); suitable 
experiments illustrating it are described on p. 26. 

To find an experimental test of the Atomic Theory's assertion that 
atoms combine in small whole numbers 

If atoms combine in small whole numbers, the compounds formed 
from elements with atoms A and B must be of the form: A,B„ 
AjB 2 , A 2 B„ A a B 3 , and so on. From this it follows that the mass A 
of one element may combine with the masses B, 2B, £B, l^B, etc. 
of the other element. But the mass A is a constant mass, because 
all atoms of a given element are identical. Similarly, the mass B is a 
constant mass. That is, the masses of B which combine with a 
constant mass of A in their various compounds should be in the 
ratio of 1 : 2 : 1 /2 : 3/2, which (multiplied by 2) is the ratio 2:4:1:3 
This is a ratio of small whole numbers. That is, if the atomic theory 
is true, the different masses of element B which combine with a fixed 
mass of element A are in a simple whole-number ratio. Experiment 
shows that these deductions are correct. The Law of Multiple Pro- 
portions expresses the general situation and is stated, with experi- 
mental illustrations, on p. 28. 

Another law, the Law of Reciprocal Proportions, can also be 
shown to follow from the portions of the Atomic Theory we are 
considering. This is dealt with on p. 30. 

We have now obtained, from the Atomic Theory, certain con- 
clusions which can be made the subject of experimental test. It is the 
magnificent achievement of Dalton to have been the first to state the 
Atomic Theory, deduce from it these conclusions suitable for ex- 
perimental check, and show that the experimental results support the 
Theory. It is a piece of work which stamps Dalton as a scientific 
genius of a very high order. 


Having reached the conclu sion that elements are made up of atoms 
scientists needed some means of denoting atoms. Dalton invented a 
system, which is now purely of historical interest and is dealt with 



briefly on p. 22. The modern simple system of representing atoms 
was suggested by Berzelius, and consists generally of using the initial 
letter of the name of the element to stand for one atom of it, e.g., one 
atom of hydrogen is denoted by H, one atom of oxygen by O, one 
atom of nitrogen by N. This rule cannot be universally applied 
because 100 elements have to share 26 letters. The difficulty has been 
readily overcome by using, for some of the elements, a symbol con- 
sisting of the initial letter, as a printed capital, together with one 
small letter from its name; for example, one atom of each of the 
elements carbon, chlorine, cerium, calcium and cjesium is denoted 
by the symbol C, CI, Ce, Ca and Cs. In the case of the metals, the 
Latin names have sometimes been used as the source of the symbol 
for example, 




























These last two metals were unknown to the Romans, but a kind of 
pseudo-Latin name has been bestowed upon each, and from this its 
symbol is taken. A list of the symbols of the known elements is given 
on p. 535. 

Definition. The symbol of an element consists of one or more letters 
which denote one atom of the element. 

It is very important to keep it clearly in mind that the symbol of 
an element does stand for a perfectly definite amount of it, and that 
amount is one atom. 


We have seen that the smallest possible particle of an element is 
called an atom. It is obvious that the smallest possible particle of a 
compound must contain at least two atoms because a compound must 
contain at least two elements and cannot contain less than one atom 
of each. To the smallest possible particle of a compound is given the 
name 'molecule'. 

The word molecule has, however, a wider meaning than this. We 
have seen that the smallest possible particle of an element is called 




one atom of it, but it does not follow, necessarily, that single atoms 
are the only particles normally existing in a mass of an element. 
Actually most elements usually exist as a mass of more complex 
particles, consisting of a number of atoms associated together and 
moving as a single particle. To this more complex particle is given 
the name 'molecule'. The distinction between the atom and the mole- 
cule of an clement is that the atom is the smallest particle of it which 
can ever be obtained, and is the unit which is concerned in chemical 
reactions, while the molecule is the smallest particle of the element 
which is normally capable of a separate existence. 

Try to get this distinction quite clear; the atom is the 
smallest particle which can participate in chemical reactions 
while the molecule is the smallest particle which can 
normally exist when the element is not concerned in 
chemical reaction. 

Definitions. The atom is the smallest, indivisible particle of an element 
which can take part in chemical change. 

The molecule of an element or compound is the smallest particle of it 
which can normally exist separately. 

Actually, most of the elementary gases consist of molecules each 
containing two atoms. The molecules of hydrogen, oxygen, nitrogen 
and chlorine are all of this type; a proof of this statement cannot be 
given until we have dealt with the Molecular Theory more fully, but 
it will be found, applied to hydrogen, on p. 99. This state of com- 
bination of the atoms is indicated by writing the molecules of these 
gases as H 2 , O,, N» and Cl«, meaning a single unit consisting of two 
atoms of each of the gases. (2H or 2CI would mean two separate 
atoms of each of these gases, a condition in which they do not 
normally exist.) The number of atoms in a molecule of an element is 
called its atomicity. 

Definition. The atomicity of an element is the number of atoms in one 
molecule of it. 

* <b 8 

Molecules of hydrogen, 
diatomic, H, 


Molecules of ozone, 
Iriaiomic, : , 

Fig. 2. 



Molecules of phos- 
phorus (in certain sol- 
vents), polyatomic, P« 

A molecule containing one atom is said to be monatomic, 
e.g., He, helium. 


- ■ i 

- *Q 


A molecule containing two atoms is said to be diatomic, 
e.g., H 2 , 2 , N 2 , Cl 2 . 

A molecule containing three atoms is said to be triatomic, 
e.g., 3 , ozone. 

A molecule containing more than three atoms is said to be 
polyatomic, e.g., P 4 phosphorus, S„ sulphur. 


From Berzelius' system of symbols is derived a simple method of 
denoting molecules of a compound or element 

Anticipating a little, we have already seen that a molecule of an 
element is denoted by writing the symbol of the element and, to the 
right and below it, a number expressing the number of atoms in the 
molecule; for example, 

H 2 denotes one molecule of hydrogen containing two 

P 4 denotes one molecule of phosphorus containing four 

S 8 denotes one molecule of sulphur containing eight 

The same device is adopted in representing the molecules of com- 
pounds, though here, of course, at least two symbols must appear 
because at least two elements must be present. Again, the small 
figure, to the right of a symbol and below it, expresses the number of 
atoms of the element present, the figure 1 being omitted. A few 
examples will make the idea clear. 

CuO denotes one molecule of copper oxide containing 
one atom of copper and one atom of oxygen. 

H 2 denotes one molecule of water containing two atoms 
of hydrogen and one atom of oxygen. 

H 2 S0 4 denotes one molecule of sulphuric acid containing 
two atoms of hydrogen, one atom of sulphur and four 
atoms of oxygen. 

CaC0 3 denotes one molecule of chalk (calcium carbon- 
ate) containing one atom of calcium, one atom of carbon 
and three atoms of oxygen. 

The close proximity of the symbols denotes that the elements are 
in chemical combination. The collection of symbols and numbers 
which together denote one molecule of a compound is called its 
formula. Thus, CuO, H a O, H 2 S0 4 and CaC0 3 are respectively the 



formula of copper oxide, water, sulphuric acid and calcium car- 

When a group of symbols is common to a class of compounds, it is 
frequently written as a bracketed group in their formula;, together 
with a number to indicate the number of groups present. For example, 
all metallic nitrates arc derived from nitric acid, HNO s , and they all 
contain the nitrate group or radical, NO s , in their formula;. When 
formula of nitrates are written, this group is preserved intact, and, 
if two or more are needed, the number is indicated by enclosing 
the N0 3 group in a bracket and writing the number needed below 
and to the right. This arrangement is convenient because it em- 
phasises the relation of the nitrates to nitric acid. For example, 
Ca(N0 3 ) 2 means the same as CaN 2 O , because the 2 multiplies 
everything inside the bracket, but Ca(NO s ) 2 indicates the relation 
of calcium nitrate to nitric acid, HN0 3 , more clearly than does 
CaN 2 8 . Similarly the formula of aluminium nitrate is written 
A1(N0 3 ) 3 rather than A1N 3 0„. 

The sulphate radical, S0 4 , which is common to sulphuric acid, 
HjS0 4 , and to all sulphates, is similarly treated. Thus the formula of 
aluminium sulphate, written as AI 2 (S0 4 ) 3 , indicates the derivation 
of this compound from sulphuric acid, H 2 S0 4 , more clearly than if 
written as A1 2 S 3 0, 2 . The hydroxyl group, OH, is also preserved in 
formulae to emphasise the relation of hydroxides to water, H.OH, 
which is regarded as hydrogen hydroxide. Thus the formula of ferric 
hydroxide is written Fe(OH) 3 , not Fe0 3 H 3 . 

If it is necessary to indicate a number of molecules of a com- 
pound, this is done by writing the appropriate number before the 
formula of the compound; for example, 

2H 2 S0 4 means two molecules of sulphuric acid. 
8HN0 3 means eight molecules of nitric acid. 
4HC1 means four molecules of hydrochloric acid. 
10H 2 O means ten molecules of water. 

It is important to notice carefully that the figure in front of the 
formula multiplies the whole of it. 2H 2 S0 4 , for example, denotes two 
molecules of sulphuric acid each containing two atoms of hydrogen, 
one atom of sulphur and four atoms of oxygen, or four atoms of 
hydrogen, two atoms of sulphur and eight atoms of oxygen in all. 

It is a common mistake of beginners to think that the 
figure multiplies only the symbol which immediately follows 
it, for example, that in 2H 2 S0 4 the 2 multiplies only the H, 
and not the S0 4 . This is quite wrong. The 2 multiplies the 
whole of the formula HgSOj. 



Notice that the formula of a compound denotes the perfectly 
definite amount of one molecule of it. 

Dalton's system of symbols 

Dalton invented a number of symbols for the atoms of elements, a 
few of which are: 







Fig. 3. 



He then indicated the formulae of compounds by combining the 
necessary numbers of these symbols, writing each one separately. 
Using modern knowledge of these compounds, his formula; would 


Methane Sulphuric acid 



# $88 



CH 4 H,SO, 
Fig. 4. 


This very laborious system is quite unsuited to the representation 
of complex molecules. Think, for example, of the task of represent- 
ing on this system the formula of cane sugar, C 12 H 22 0,,. It would 
entail the drawing of 12 separate symbols for carbon, 22 for hydrogen 
and 1 1 for oxygen, and the result would be an unwieldy and confusing 
jumble of 45 separate signs. The system was soon abandoned for that 
of Berzelius. 


It is now possible for us to represent chemical reactions by means 
of statements involving only formula; and symbols. 

For example, if we wish to represent the fact that one molecule of 
copper oxide reacts with one molecule of sulphuric acid, producing 
one molecule of copper sulphate and one molecule of water, we can 
do so in the form: 

CuO + H a S0 4 -> CuS0 4 + H 2 

This kind of statement is called 'a chemical equation'. 
The + sign on the left of the equation means 'reacts with', but on 
the right it means simply 'and*, while the arrow means 'producing'. 

questions 23 

Take another simple chemical equation. 

Zn + H,S0 4 ->- ZnS0 4 + H 2 

This means: 'one atom of zinc reacts with one molecule of sulphuric 
acid producing one molecule of zinc sulphate and one molecule of 

Again, the formula of potassium chlorate being KCI0 3 , and of 
potassium chloride KC1, the equation 

2KC10 3 -> 2KC1 + 30 8 

means: 'two molecules of potassium chlorate decompose producing 
two molecules of potassium chloride and three molecules of oxygen'. 
It is important to notice that there must be the same number of 
each kind of atom on the right of a chemical equation as on the left. 
Otherwise the equation would imply that atoms had been created or 
destroyed, which is impossible. This process of equalisation is called 
'balancing' (see p. 58). 


1. State the meaning of the following formula; : KNO a (nitre), 
CuSO,.5H,0 (blue vitriol), PbCl s , NajCO s .10H.O (washing soda), e.g., 
the formula H 2 (water) means that one molecule of water contains two 
atoms of hydrogen and one atom of oxygen. 

2. How many atoms of the various elements arc indicated by the 
following formulas: 

2H s O, 5HCI, 7HN0 3 , 20PbSO 4 , IICu(NO I ) 1 ? 

3. Explain why the terms 'atom* and 'molecule' can both be correctly 
applied to particles of an clement, but only one of them to particles of a 

4. What do you understand by the term 'atomicity'? What is the 
atomicity usually assigned to the elements helium, chlorine, ozone, 

5. Why is the formula for lead nitrate written Pb(N0 3 ) a ? (It could also 
be written PbN,0«.) Give three similar examples. 

6. Using the Atomic Theory as your illustration, show the importance of 
the balance in determining the truth of a proposed theory. 

7. State briefly Dalton's Atomic Theory. Explain why the theory is 
named after him in spite of the fact that he was not the first to bring forward 
the idea that matter consists of atoms. 

8. State the Law of Multiple Proportions and show how it is explained 
by the Atomic Theory. 

An element X forms two oxides containing 77.47 and 69.62 per cent of 
X respectively. If the first oxide has the formula XO, what is the formula 
of the second oxide? (C.) 

Chapter 3 

Laws of Combination by Weight (experimental) 

The Law of Conservation of Mass 

THE Law of Conservation of Mass (or Indestructibility of Matter) 
states that matter is neither created nor destroyed in the course 
of chemical action. 
Let us illustrate this by an experiment. 

Into a conical flask put some silver nitrate solution and lower into 
it carefully by means of a thread a small test-tube full of hydro- 
chloric acid. Insert the stopper (Fig. 5). 
Place the flask on the pan of the balance 
and weigh it. 

Note that you have just weighed 
(besides those portions of the 
apparatus which are unchanged 
throughout) some water and silver 
nitrate and hydrogen chloride. 

Allow the two liquids to mix by tilting 
the flask a little, and you will observe a 
white precipitate of silver chloride whilst 
nitric acid will be formed in solution. 

AgN0 3 + HC1 ->- AgCl + HN0 3 

Replace the flask on the pan of the balance 
and weigh again. You will find the weight 
is exactly the same as before. 

' acid 




Fig. 5. 

Experiment to illustrate 

the Law of Conservation 

of Mass. 

This time you were weighing (besides those portions of 
the apparatus which are unchanged throughout) some water 
and silver chloride and hydrogen nitrate (or nitric acid). 

Hence, although substances may undergo chemical changes, the 
total mass of the products of the reaction is exactly equal to the 
total mass of the reacting substances. 




Other substances suitable for the above experiment are given in 
the following equations, which illustrate the chemical changes: 


2NaN0 3 

BaCl 2 + H 2 S0 1 ->BaS0 4 
Pb(N0 3 ) 2 + 2NaCI -»- PbCl a 

There are many other examples. 

These experiments are concerned with solutions. This is merely a 
matter of experimental convenience. It is difficult to weigh gases, 
and solids are unsuitable because they do not generally undergo rapid 
and complete reactions. The same experimental results are obtained, 
however, when solids and gases are involved. 

It must be noted, however, that a law of the kind we are dis- 
cussing here is valid only within the limits of experimental error in- 
volved. A very accurate set of experiments carried out by Landolt 
about 1906 showed that the Law of Conservation of Mass is valid 
(for the cases investigated) to about one part in ten million. For 
most purposes, errors of this order are negligible. It should be noted 
also that, according to modern ideas, emission of energy during 
chemical change involves loss of mass. The relation is expressed in 
Einstein's equation, E = mc-, where E is the energy, m the mass and 
c the velocity of light (in appropriate units). In ordinary chemical 
changes, the accompanying energy changes are so small that their 
influence on mass cannot be detected by any weighing apparatus in 
common use. Nevertheless, the true conservation is one of mass-and- 
energy together. 

Apparent destruction of coal 

The above experiment may not seem conclusive to you because in 
it there seems little possibility of loss. If you consider the burning of 
coal, where only a small ash is left, it seems much more likely that the 
matter of the coal has been destroyed. The only real difference is that 
some of the reactants and products are invisible gases. Actually there 
is no loss of matter at all. If we could weigh all the oxygen which 




00 #+(coal} 





Weight of this = 

8>8> <£> 
8>8>«P + <V > o> + 

Small amount 
of soot, ash 
and other 


Weight of this 

Fig. 6. 



burns the coal, and all the ash, soot, water-vapour and carbon dioxide 
into which the coal is changed by the burning, we should again find 
that the total mass of the materials before the reaction was the same 
as the total mass of the total products after it (Fig. 6). 

The Law of Definite Proportions 
(or Constant Composition) 

The Law of Definite Proportions states : All pure samples of the 
same chemical compound contain the same elements combined in the 
same proportions by mass. 

The illustration of this law entails the performance of several 

Principle of the Experiment. Copper(II) oxide (i.e., cupric oxide) 
is prepared by several different methods and the samples are analysed 
by reduction in a stream of hydrogen, or coal-gas, and shown to 
contain copper and oxygen in the same proportions. 

Preparation of the samples of copper oxide (copper(II) oxide) 

Sample A. Starting from Copper. Place a little clean, pure copper 
foil in a large crucible in a fume-chamber and carefully add concen- 
trated nitric acid a little at a time. Brown fumes of nitrogen dioxide 
(poisonous) are seen and green copper(II) nitrate solution is formed. 

Cu + 4HN0 3 -► Cu(N0 3 ) 2 + 2H 2 + 2NO, 

Evaporate the solution to dryness, then heat the green solid copper(II) 
nitrate until no more brown fumes of nitrogen dioxide are evolved. 
The black solid left is the first sample of copper(II) oxide. Store it 
in a desiccator to keep it dry. 

2Cu(N0 3 ). -»• 2CuO + 4N0 2 -f 2 
Sample B. Starting from Copper(II) Sulphate. Put some copper(H) 
sulphate solution into a beaker and add excess of caustic soda solu- 
tion. A blue gelatinous precipitate of copper(Il) hydroxide appears. 
Heat the beaker and its contents on a tripod and gauze by means of a 
Bunsen burner. The precipitate changes to black copper(II) oxide. 

CuS0 4 + 2NaOH -► Cu(OH) 2 + Na t S0 4 

Cu(OH) 2 -> CuO + H a O 

Filter off the black solid, wash it several times with hot distilled 
water, and allow it to dry in a hot oven or on a porous plate. 

Transfer the oxide to a crucible and heat it with a burner to drive 
off the last traces of water. Store the oxide in a desiccator. 

Sample C. Starting from Copper(If) Carbonate. Place a little 
copper(II) carbonate in a dry crucible and warm it gently. It decom- 



poses, turning from green to black, and carbon dioxide is given off. 
(Test — a drop of lime-water on the end of a glass rod is turned milky. 
The milkiness is caused by a precipitate of chalk.) Black copper(II) 
oxide is left. After heating it for some time, put the oxide into a 
desiccator to keep it dry. 

CuC0 3 — >- CuO + C0 2 

Analysis of the samples of copper(TI) oxide by converting them to 
copper by heating in coal-gas. 

Weigh three clean, dry porcelain boats. Put into the boats 1-1$ 
grams of the three samples of copper(II) oxide. Weigh all the boats 
again. Put them into a hard glass tube and connect up the apparatus 


Fig. 7. 
Reduction of three samples of copper(II) oxide by coal-gas. 

as shown in Fig. 7. The tube must slope so that the end C is the 

Turn on the coal-gas and light it at the jet D. Then heat each boat 
in turn. All the samples of the oxide will glow and leave reddish- 
brown copper. The water formed by the combination of the oxygen 
of the oxide with the hydrogen of the coal-gas condenses at C, where 
the tube is cooler. The end C is lower than the end A to prevent this 
water from running back on to the hot part of the tube, which might 
be broken. When the action is complete, allow the tube to cool, 
keeping the coal-gas stream passing so that air cannot enter and 
oxidise the copper again. When the boats are cool, weigh all the 
three boats again. Work out the results as shown below. 

CuO + H 2 ^ 


Cu + H 2 

Within the limits of experimental error, the percentages of copper 
(and hence of oxygen) in all three samples of copper oxide are the 



same. It is found that, however samples of a given compound are 
prepared, they always contain the same elements in the same pro- 
portions by mass. 

Specimen Results 

Sample A 

Sample B 

Sample C 

Weight of porcelain boat 

Weight of porcelain boat and copper 

Weight of porcelain boat and copper 
Weight of copper 
Weight of copper oxide 

Percentage of copper in copper oxide 

3.01 g 

4.26 „ 

4.02 „ 
1.01 „ 

1.01 x 100 


= 80.8 

2.50 g 

3.65 „ 
3-42 „ 
0.92 „ 
0.92 x 100 
= 80.0 

2.70 g 

3.85 „ 
1.15 x 100 
= 79.8 

The Law of Multiple Proportions 

The Law of Multiple Proportions states: If two elements A and B 
combine together to form more than one compound, then the several 
masses of A, which separately combine with a fixed mass of B, are in 
a simple ratio. 

Coppcrfl) oxide 

Copper(Il) oxide 

Weight of boat 

6.90 g 

7.30 g 

Weight of boat and oxide 

9.75 „ 


Weight of boat and copper 



.•. Weight of copper 

2.53 „ 


Weight of oxygen 


0.38 „ 

.'. 0.32 g oxygen 

.". 0.38 g oxygen 

is combined with 

is combined with 

2.53 g copper 

1.52 g copper 

100 g oxygen are 

100 g oxygen are 

combined with 

combined with 

2.53 x 100 

1.52 x 100 

0.32 gcoppcr 

0.38 BCOpper 

= 790g 

= 400g 

This ratio is 2 : 1 withii 

l limits of experimental 

Principle of Experiment. Copper and oxygen form two oxides, 
cupric oxide (copper(II) oxide, CuO) and cuprous oxide (copper(I) 
oxide, Cu 2 0). Pure samples of these two are reduced in a current of 



hydrogen and the masses of copper which combine separately with, 
say, 100 grams of oxygen in the two compounds, are calculated from 
the weighings. 

Weigh two clean dry porcelain boats and weigh them again con- 
taining samples of pure, dry, copper(I) oxide and copper(ll) oxide 
respectively. Reduce the oxides to copper in a stream of dry hydro- 
gen, or coal-gas, as described in the last experiment. The preceding 
table is a list of weighings and a specimen analysis. 

Hence the masses of copper which have separately combined 
(i.e., to form the two different oxides) with a fixed mass, 100 g, of 
oxygen are in the ratio 2 : 1. The law could have been illustrated 
just as easily by fixing the mass of copper. 

If desired, the results of the above experiment can be expressed in 
the following way, given the atomic weights of oxygen and copper as 
16 and 63.5 respectively. (For mole, see p. 39.) 

In copper(I) oxide 

0.32 g of oxygen combine with 2.53 g of copper, so -^—- moles of 


2 53 
oxygen atoms combine with --— moles of copper atoms. That is, 


0.02 moles of oxygen atoms combine with 0.04 moles of copper 

atoms, i.e., the ratio of moles is 1 : 2. 

In copper(II) oxide 


0.38 g of oxygen combine with 1.52 g of copper, so -f— moles of 


oxygen atoms combine with -^— moles of copper atoms. That is, 


0.024 moles of oxygen atoms combine with 0.024 moles of copper 

atoms, i.e., this ratio of moles is 1 : 1. From this, the numbers of 

moles of copper atoms which combine with one mole of oxygen 

atoms in the two compounds are in a simple ratio to one another 

(2 : 1). Since one mole of oxygen atoms is a fixed mass, this result 

agrees with the Law of Multiple Proportions. Also, the figures for 

mole ratios point to Cu 2 and CuO as the simplest formulae for 

copper(I) oxide and copper{II) oxide respectively. They were formerly 

known as cuprous oxide, Cu 2 0, the red oxide, and cupric oxide, 

CuO, the black oxide. 

Analysis of mercurous and mercuric chlorides as an illustration of the 
Law of Multiple Proportions 

The accuracy of the experiment described above depends on the 
purity of cuprous oxide, a material extremely difficult to obtain in a 



pure state. Mercurous and mercuric chlorides can both be obtained 
in a high degree of purity (Analar quality is used), and the following 
experiment may be performed as a class illustration of the law. 
Weigh a boiling-tube, add two or three grams of mercurous chloride, 
and weigh again. Add a teaspoonful of sodium hypophosphite (see 
p. 458), half fill with water, immerse in a beaker of water, and warm. 
Repeat using mercuric chloride, taking care to distinguish between 
the two boiling-tubes. After about 20 minutes, globules of mercury 
are seen in each tube. Wash the mercury by decantation several times 
with water (pouring the water into a beaker so that if mercury is lost, 
it may be retrieved), then with methylated spirit, and finally with 
ether. (Care! Extinguish all flames in the vicinity.) Replace the tubes 
in the warm water for a minute to remove traces of ether, dry the 
outside, and weigh. Calculate the separate masses of chlorine associ- 
ated with, say, one gram of mercury in each of the two chlorides. 
These masses will be in the ratio 1 : 2. 

NaHjPO, + 2HgCl 2 + 2H a O 



(powerful reducing 


■ NaH 2 PO« + 4HC1 + 2Hg 

acid sodium 

The Law of Reciprocal Proportions 

This is a fourth law which can be deduced from the Atomic 
Theory. It is expressed in the statement: 

If an element A combines with several other elements, B, C, D, the 
masses of B, C, D, which combine with a fixed mass of A are the 
masses of B, C and D which combine with each other, or simple 
multiples of those masses. 

Like the Law of Multiple Proportions, this law can be derived 
from Dalton's assumption that combination takes place between 
small whole numbers of atoms, all the atoms of a particular clement 
being identical. 

Relation between scientific law and theory 

A scientific law is simply a generalised statement or observed facts. For 
example in ail cases examined, it has been found that the total mass of 
the products of a chemical action is equal to the total mass of the original 
reagents. This result is assumed to apply to <7//such cases and is generalised 
as th £., k? w °l Conservation of Mass. A scientific law is subject to two 
possibilities of error; first, it is not possible to examine every last case of 
the operation of the law. There may be millions of them. Second, the law 
can be accurate only to the limit of experimental error. 
i,*™™' IS u an ldea P ut forward to explain the existence of one or more 
laws. In the above case, the law is explained by a part of Dalton's atomic 



theory, i.e., the idea that atoms exist which can be neither created nor 
destroyed in chemical action. Once a scientific law has been definitely 
established it will usually stand for any relevant time, being well grounded 
in fact. A theory put forward to explain the law may, however, be replaced 
from time to time as knowledge increases. 


1 . 1 g of one oxide of X contained 0.5 g of X, and 4 g of another oxide 
of X contained 1 .6 g of X. Show these weights to be in accordance with 
the Law of Multiple Proportions. 

2. Three oxides of nitrogen contain 53.3%, 69.6% and 36.4% of oxygen 
respectively. Show that the Law of Multiple Proportions is upheld. 

3. Show that the following results obtained by the reduction of the two 
oxides of a metal are in agreement with the Law of Multiple Proportions. 

Weight or boat 
Weight of boat + oxide 
Weight of boat + metal 

1st Compound 

2nd Compound 

5.30 g 
13.85 „ 

4.45 g 
13.05 „ 
12.08 „ 

4. 1.90 g of one oxide of copper gave 1.52 g of copper on reduction. 
2.85 g of another oxide gave 2.53 g of copper on reduction. Show these 
results to be in accordance with the Law of Multiple Proportions. 

5. Two oxides of a metal contained respectively 7.41% and 3.85% of 
oxygen. Show these facts agree with the Law of Multiple Proportions. 

6. The sulphides of a certain element contained 33.7% and 20.4% of 
sulphur. Do these figures agree with the Law of Multiple Proportions? 

7. State the Law of Multiple Proportions. Describe how you would 
attempt to verify this law experimentally, if you were given specimens of 
litharge (PbO) and lead dioxide (Pb0 2 ). Name one other pair of substances 
which you could use to verify this law. (N.U.J.B.) 

8. A metal forms two oxides. 1.000 g of each oxide contains 0.239 and 
0.385 g of oxygen respectively. Determine the equivalent of the metal in 
each oxide and show that these figures are in agreement with the Law of 
Multiple Proportions. (C.) 

9. State the Law of Constant Composition (that is, Definite Proportions). 
You are required to verify this law by using the black oxide of copper, and 
for the purpose you are supplied with specimens of copper nitrate and 
copper carbonate. State clearly (a) how you would prepare specimens of 
copper oxide; (h) how you would use them to verify the law. (N.U.J.B.) 

10. State the Law of Definite Proportions. Describe fully the experi- 
ments you would carry out in the laboratory in order to prove the truth of 
the law in the case of one chemical compound. (D.) 

1 1 . A metal, X, forms two oxides, A and B. 3.000 g of A and B contain 
0.720 g and 1.160 g of oxygen respectively. Calculate the masses of metal 
in grams which combine with one gram-atom of oxygen in each case. What 
chemical law do these masses of metal illustrate? Explain briefly. If the 
oxide B has the formula XO, what is the formula of oxide A? (O = 16.) 



12. Three samples. A, B and C, all known to be oxides of a metal, M, 
gave, by reduction in hydrogen, the following figures: 1.60 g of A left 
1.28 g of M; 1.44 g of B left 1.28 g of M; 1.00 g of C left 0.80 g of M. 
Calculate the masses of M that combine with one gram-atom of oxygen 
in each case. Explain briefly what two chemical Laws are illustrated by the 
results? If oxide A has the formula MO, what are the formula; of B and C? 
(O = 16.) 

1 3. When heated in dry chlorine, 1 . 1 2 g of a metal, M, were converted 
to 3.25 g of anhydrous chloride, A. When heated in dry hydrogen chloride, 
1 .400 g of M were converted to 3. 1 75 g of anhydrous chloride, B. Calculate 
the masses of M which combine with one gram-atom of chlorine in these 
chlorides. What chemical Law do these figures illustrate? Explain briefly. If 
the formula of A is MC1 S , write formula: for B and for the corresponding 
sulphate of M. (CI = 35.5.) 

14. Two oxides, A and B, of a metal, M, contained 70.0% and 77.8% 
respectively of the metal. Calculate the masses of M that combine with 
one gram-atom of oxygen in A and B. What chemical Law do these 
figures illustrate? If B corresponds to a chloride MCl a , what is the formula 
of the chloride corresponding to A? (O = 16.) 

Chapter 4 

Atomic Weight 

Introduction of the chemical balance 

AS soon as the existence of atoms was recognised, it was obviously 
desirable to try to obtain as much information as possible about 
the atoms of different elements, and, more particularly, to show how 
their weights compared with one another. This was necessary to 
enable quantitative experiments of various kinds to be made, to test 
doubtful points about the Atomic Theory, and also to enable 
chemists to make calculations of the quantities of materials involved 
in their experiments. Indeed, until the work of chemists could be 
made the subject of accurate weighings, and calculations could be 
tested by experiment, little progress could be expected. This was 
recognised more particularly by Black, a Scottish scientist working 
in Glasgow, and his name will always be honoured among chemists 
for his persistent and pioneer use of the chemical balance for checking 
his ideas. He taught scientists the importance of obtaining definite 
quantitative results instead of the vague qualitative statements with 
which they had been satisfied. The question for chemists became not 
only 'What happens?' but also 'What weight of each material is 
involved when it happens?' 

Chaos of 1820-1850 

We must not hide from ourselves the fact that the project of finding 
out how the weights of different atoms compared with each other 
proved extremely difficult to carry out. For about forty years after 
the Atomic Theory was suggested by Dalton, Chemistry was in a state 
of chaos. It is difficult for us now to read with understanding any 
book on Chemistry written between about 1820 and 1850, because 
chemists were simply groping about trying to solve the problem of 
comparing the weights of atoms, and they were making little progress. 
Let us try to understand where the difficulty lay and to follow the 
stages by which full knowledge was finally achieved. 

In the first place, we must understand that chemists were not 




attempting to obtain the weights of individual atoms. It was clearly 
recognised that atoms were very small indeed and that there was no 
hope whatever at that time of obtaining the actual weight of a single 
atom of any element. The question was, rather: 'How do the weights 
of the atoms of different elements compare with one another ? Is, for 
example, a sodium atom heavier than an atom of oxygen, and, if so, 
how many times? Is a silver atom heavier than an atom of gold, and, 
if so, how many times?' Actually, we can today find the real weight 
in grams of a single atom of any element; we have accomplished what 
the chemists of 1830 regarded as a vain aspiration, but the process 
has taken 100 years, and when you read that the weight of a hydrogen 
atom is 0.000 000 000 000 000 000 000 0017 g, you will not be sur- 
prised that science has consumed a great deal of time, and the patient 
efforts of many, to reach a stage at which such a minute quantity can 
be reasonably accurately measured. 

Meaning of the term atomic weight 

Returning to the simpler problem of comparing the weights of 
atoms, it became necessary first to fix some standard of weight with 
which all the atoms could be compared. We were measuring potatoes 
in stones, jam in pounds, and masses of steel in tons. In what weight- 
units could we express the weights of atoms? To compare weights 
of atoms by using grams would, in fact, have been far less sensible 
than trying to express the weight of grains of sand in tons. Chemists 
recognised this and decided to compare the weights of all other 
atoms with the weight of a hydrogen atom. 

The hydrogen atom was chosen because it is the lightest 
of all the atoms and would give numbers greater than 
unity for the comparative weights of all the other heavier 
atoms. Tilts is more convenient than working in fractions or 
decimals, which the choice of any other atom as standard 
would necessitate. 

Definition. The atomic weight of an element is the number of times 
one atom of the element is as heavy as one atom of hydrogen. 1 

You must understand clearly what this most important character- 
istic of an atom is. If you look in a 'Table of Atomic Weights', you 
will see some statement such as O = 16 or Na = 23 or P = 31. This 
is the chemist's shorthand way of saying 'The atomic weight of 
oxygen is 16, of sodium 23, and of phosphorus 31'; or, more fully, 
'Every atom of oxygen is 16 times as heavy as any atom of hydrogen'; 
or 'Every atom of sodium is 23 times as heavy as any atom of 

1 For modernisation of this definition, sec pp. 37, 39. 



hydrogen'; or 'Every atom of phosphorus is 31 times as heavy as any 
atom of hydrogen'. All this is conveyed by O -= 16, Na = 23, P = 31. 
You do not yet know how these figures have been obtained; take 
them for granted for the moment. It will be obvious that if the 
weights of all atoms are compared with the weight of a hydrogen 
atom in this way, the weights of all atoms are also compared with 
each other. When we say, for example, that the atomic weight of 
oxygen is 16 and that of sodium is 23, we also state that the sodium 
atom is heavier than the oxygen atom in the proportion of 23 to 16. 

The fact that we have chosen the weight of a hydrogen atom as our 
unit of atomic weight leads to the statement that the atomic weight 
of hydrogen is 1 (H — 1). 

It is most important to get this idea of comparing the weights of 
all other atoms with that of a hydrogen atom firmly fixed in your 
mind. Consider it for a moment from another angle. A fairly com- 
plicated set of chemical experiments and theoretical arguments has 
established the conclusion that the atomic weight of sulphur is 
32 (S =■ 32). In the long run, this simply means that if we could 
separate out from a piece of sulphur one single atom of it, place it on a 
sufficiently delicate see-saw and then, to the other side of the see-saw 

ATOM (S- 32) 

weighs k much *j 



Fig. 8. 

add, one at a time, some hydrogen atoms, the see-saw would just 
balance when exactly 32 hydrogen atoms had been added (Fig. 8). 

In actual practice, we cannot separate out the one single atom of 
sulphur and the single atoms of hydrogen in this way, but, if we 
could, all the chemical experiments which have been required to find 
the result S = 32 would be unnecessary. 

Molecular weight 

The weights of molecules of elements or compounds are also com- 
pared with the weight of a hydrogen atom. 

Definition. The molecular weight of an element or compound is the 
number of times one molecule of it is as heavy as ONE ATOM of 
hydrogen. 1 

1 For modernisation of this definition, see pp. 37, 39. 



Early method for determining atomic weights 

Considerable confusion existed in the chemical world in the first 
half of the nineteenth century because of the difficulty of distinguish- 
ing between atomic weights of elements and their combining weights. 
For example, if two elements, A and B, combined in equal numbers 
of atoms to give a compound of simplest formula AB, then their 
combining weights were also their atomic weights. If, however, the 
compound formed was A 2 B, the combining weight of A (in relation 
to the atomic weight of B) was twice its atomic weight. 

These difficulties were eventually overcome by the evolution of a 
system in which the combining weight (called the equivalent) of an 
element was referred to one gram of hydrogen and its combining 
power (or valency) was referred to one atom of hydrogen. This led 
to the definitions which follow. 

The gram equivalent weight of an element is that number 
of grams of the element which combine with, or displace, 
one gram of hydrogen. 

The valency of an element is the number of hydrogen 
atoms which combine with, or displace, one atom of the 

From these definitions, it can be shown that the following relation 
exists for any element: 

Equivalent x Valency = Atomic weight 

In this relation, atomic weight is & fixed quantity. The valency of an 
element may, however, vary while always retaining a whole-number 
value. It will be shown later that this can occur because combining 
power is essentially an electronic phenomenon and an atom may 
vary its use of electrons, and the number used for combining 
purposes, according to its environment (within certain limits). This 
means that an element may have more than one equivalent. For 
example, iron may use two electrons per atom for valency purposes 
(and is then known as iron{II)) and its equivalent (28) is half its 
atomic weight (56); or it may use, as iron(HF), three electrons per 
atom and have an equivalent of 18.7, i.e., one-third of its atomic 

Difficulties and uncertainties of this sort were gradually overcome 
in the second half of the nineteenth century, and, before 1900, the 
atomic weights of about 90 elements were known with substantial 
accuracy. At this time, the atomic weight scale was still based on 
H = I, which (from the composition of water) leads to O = 15.88. 
Early in the twentieth century, it was agreed that the scale should 
be modified slightly by basing it on O = 16.00, which raised all 



atomic weights by something under 1%, making, for example, 
H = 1.008 (instead of exactly 1). The scale then remained unchanged 
until 1962 when (for reasons connected with the mass spectrometer), 
it was modified very slightly (a change of only 37 parts per million) by 
basing it on the carbon isotope, £Ci which was given the atomic 
weight of exactly 12 units. The older methods of atomic weight 
determination have now been substituted by mass spectrometry. A 
brief account of the mass spectrometer is given below. 

With the present 12 C-scale of atomic weight, the gram-equivalent 
of an element must theoretically be related to 3 g of carbon, since the 
valency of carbon is 4. For practical purposes, however, the cor- 
responding weights of elements such as hydrogen, oxygen and chlor- 
ine are much more important and a working definition of equivalent 
is the following: 

The gram-equivalent weight of an element is the number 
of grams of the element that combine with, or displace, 
3 grams of carbon-12, 1.008 grams of hydrogen, 8.00 grams 
of oxygen or 35.46 grams of chlorine. 

Experimental work relating to two or three methods of determining 
equivalents is given in Chapter 6, with some corresponding practice 
calculations at the end of the Chapter. 

Note on the equation: Equivalent X Valency = Atomic weight 

Recently, there has been objection to the use, by text-books, of this 
equation on the ground that it does not (superficially) apply correctly 
to certain cases. In particular, it is objected that the apparent 
equivalent of iron (21) in the oxide, Fe 3 4 , bears no simple (valency) 
relation to the atomic weight of iron (56). It is known, however, that 
the reaction of this oxide with warm, concentrated hydrochloric acid 
produces a mixture of iron(II) and iron(III) chlorides. 

Fe 3 4 + 8HCI -► FeCl 8 + 2FeCl 3 + 4H s O 

It seems, therefore, that iron, in Fe 3 4 , is acting with valencies of 2 
and 3 in the same compound, i.e., as FeO.Fe 2 3 . 

As iron(ll) or ferrous iron, its equivalent is 28, and as iron(III) 
or ferric iron, 18.7. Iron(II) is combined with one-quarter of the 
oxygen in Fe 3 4 and contributes 7 units to the apparent equivalent; 
iron(III) is combined with three-quarters of the oxygen in Fe 3 4 and 
contributes 14 units (i.e., three-quarters of 18.7). That is, the equa- 
tion, E X V = At. wt, applies accurately to the behaviour of iron(II) 
and iron(III) separately and to their joint behaviour if correctly 
apportioned. A similar case occurs in triplumbic tetroxide, Pb 3 4 , 
which acts as PbO a .2PbO. 



Modern atomic weight measurement. Mass spectrometer 

A direct mode of determination of atomic weights has been avail- 
able since 1920 through the mass spectrometer introduced by Aston. 
It must be remembered, however, that, by that date, the atomic 
weights of about 90 elements were fairly accurately known, and, at 
first, the results from Aston's work had to be checked against known 
atomic weights previously obtained from chemical experimentation. 
Since 1920, however, the mass spectrometer has been greatly im- 
proved and it can now give reliable, very accurate results. To a large 
extent, it has rendered chemical determination of atomic weights 

In principle, the mass spectrometer first produces positive ions of 
the element under investigation by causing its atoms to lose electrons. 
This may be done, for example, by very strong electrical heating of 

S, S 2 


F + 





Flo. 9. 

Mass spectrometer. 

a trace of a salt of the element on a tungsten filament, or by subjecting 
a sample of the element in gaseous form to electron bombardment. 
For a univalent element, X, 

X-vX+-f e~ 

The stream of ions, X + , is selected by slits, S x and S 2 , into a narrow 
beam which is deflected and dispersed by the electrical field between 
metallic plates at F. An electromagnetic field, M, at right angles to 
the field, F, then focuses the beam on to a photographic plate, P, so 
that all particles of the same charge/mass ratio fall on to a line 
resembling an optical spectrum line. This line is an image of the 
selecting slits. The whole process occurs at very low pressure. By 
adding oxygen, O = 16, to the element under test, early experi- 
menters could supply known lines at masses 8, 16 and 32 (represent- 
ing ions O s+ , + and O g + ) and use these as measuring standards 
for the lighter elements. Modern instruments can be made to record 
a graph on a moving paper roll. The graph shows a peak correspond- 
ing to each ion beam, from which the mass of the corresponding 



atom can be calculated and the height of the peak indicates the 
intensity of the beam. 

Aston's work of 1920 was notable not so much for its measurement 
of atomic weights as for its confirmation of the existence of isotopes. 
For example, chlorine (chemical atomic weight 35.5) showed two 
kinds of atoms (isotopes) to be present in it by producing two mass 
spectrum lines at 35 and 37 with intensities in the proportion of 
about 3 : 1. 

Notice that atomic weight itself has no units; it is simply a number 
which states the relative mass of the atom concerned on a scale which 
gives the standard atom, '„ 2 C, a mass of 12. If, however, the gram unit 
is attached to the atomic weight, the quantity of the element thus 
indicated is called one gram-atom or one mole of atoms. A mole of 
atoms of any element always contains the same number of atoms, 
6.02 x 10 23 . This figure is the Avogadro Constant, often called the 
Avogadro number; its derivation is given on p. 107. The volume 
occupied by a mole of atoms of an element may vary enormously, 
e.g., a mole of hydrogen atoms (1.008 g) occupies 1 1 200 cm 3 at s.t.p., 
while a mole of platinum atoms (195 g) occupies only 9.1 cm 3 in the 
same conditions. 

Chapter 5 

The Behaviour of Gases under Temperature 
and Pressure Change 

The nature of gas pressure 

ANY gas consists of a collection of molecules of a particular kind 
which are in a state of rapid motion. The fact that the molecules 
are in motion is evident from the fact that if a small quantity of an 
odorous gas, such as hydrogen sulphide, is liberated at any point in a 
laboratory the smell of the gas soon pervades the whole room. 

If the gas is confined in a closed vessel, some of the moving mole- 
cules strike the sides of the vessel and each impact exerts a small force 
upon the side. The number of molecules of gas inside such a vessel 
will normally be very large and, on the average, the same number of 
molecules will strike a given area on the sides of the vessel each 
second, so producing a steady pressure. 

Relation between pressure and volume 

Now suppose that one of the sides of a cylindrical vessel is a smooth 
piston and that there is a pressure, P, exerted on the piston just great 
enough to resist the pressure of the gas, of which the volume is V 

(Fig. 10, 1). The piston will remain still. 
Now suppose that the pressure on 
the piston is suddenly reduced to \P, 
without temperature change. The gas 
pressure is the greater and the piston 
will move up. As it does so the gas will 
fill the greater volume now available. The 
molecules will be more loosely packed 
in this larger space and so fewer will 
strike the sides of the vessel in a given 
time; that is, the pressure of the gas 
falls as the piston slides upwards. A 
stage will be reached when the gas 

Fia. 10. 

Effect of decrease in pressure 
on a given mass of gas. 




occupies so large a volume that its pressure has also been reduced 
to iP and the piston will then stop (Fig. 10, II). This will happen 
when the volume of the gas has doubled; that is, a halving of the 
pressure causes the volume of the gas to be doubled. Similarly, it 
would be found that if the pressure on the piston was reduced to 
^P, it would come to a stop when the volume of the gas had in- 
creased to four times its original value. 

Expressing the result generally, the pressure of a gas decreases in 
the same proportion as its volume increases. 

From this, it is clear that if we multiply the varying 
volumes of a given mass of gas by the corresponding pres- 
sures, any decrease in the value of one of them will be 
exactly counterbalanced by the increase in the value of the 
other, and the result will always be the same. 

Expressed mathematically, this may be stated in the form: 
p x v x = p t v t (temperature constant) 
where p, and p t arc two pressures and v 1 and v 2 the corresponding 
volumes of a given mass of gas. 

This result is known by the name of its discoverer as Boyle's Law. 
Boyle's Law. The volume of a given mass of gas is inversely propor- 
tional 1 to its pressure, if the temperature remains constant. 

Temperature change 

It is common knowledge that a rise of temperature causes objects 
to expand and a fall of temperature causes contraction. The rule 
applies to gases, liquids and solids but the effect is much more marked 
in the case of gases than in the case of the other two. Charles found 
that, if pressure is constant, the volume of a gas increases or decreases 
by ^ of its volume at 0°C for every °C rise or fall of temperature; 
that is, if we take 273 cm 3 of any gas at 0°C, its volume will rise or 
fall by 1 cm 3 for every °C rise or fall of temperature. Thus, at — 1 C C, 
the volume will be 272 cm 3 ; at -2°C, 271 cm 3 ; at -3°C, 270 cm 3 and 
so on. This leads to the absurdity that, if the temperature falls to 
— 273°C, the volume of the gas will be cm 3 — the gas will have 
vanished ! In actual practice, no substance can remain gaseous at such 
low temperatures; all become solids and the Law of Charles does not 
then apply, but this temperature, at which the volume of a gas would 
theoretically be reduced to zero, gives us the lowest possible tem- 
perature that can ever be reached. It is called absolute zero. 

A temperature scale is in use starting from this absolute zero as 0°, 

1 'Inversely proportional' is the mathematical expression of the fact that as 
the pressure increases the volume decreases in the same proportion. 



and using Centigrade (Celsius) degrees. This is called the Kelvin (K) 
Scale in honour of Lord Kelvin who first suggested it. Measurements 
on this scale are stated in units, K, and a general kelvin scale tem- 
perature is represented by the capital letter symbol, T. Since absolute 
zero is the same as — 273°C, it is clear that the kelvin scale starts 
measuring temperature from a point 273°C lower than the starting 
point of the Centigrade (Celsius) scale and, to convert centigrade 
temperatures to kelvin temperatures, we must add 273°. Thus: 

-253°C is the same as 20 K (-253 + 273) 

0°C is the same as 273 K (0 + 273) 

15°C is the same as 288 K (15 + 273) 

Restating the Law of Charles using kelvin temperature, we find that 
273 cm 3 of gas at 273 K (0°C) will become 274 cm 3 at 274 K (1°C), 
275 cm 3 at 275 K (2°C), 276 cm 3 at 276 K (3*C) and so on; for failing 
temperatures, the volume of gas will be 272 cm 3 at 272 K., 271 cm 3 
at 271 K, 270 cm 3 at 270 K. and so on. This gives us the rule (known 
as Charles's Law) that the volume of a given mass of gas increases 
in the same proportion as its kelvin temperature, if pressure is 

Charles's Law. The volume of a given mass of gas is directly propor- 
tional 1 to its kelvin (absolute) temperature if pressure is constant. From 
this it follows that if we divide the varying volumes of a given mass 
of gas by the corresponding kelvin temperatures, any increase in the 
volume will be exactly cancelled by the increase in the temperature, 
and the result will always be the same. This can be expressed in the 

(pressure constant) 

where u, and v 2 are the volumes of the gas at kelvin temperatures 
7", and T t respectively. 

Combining this with the equation expressing Boyle's Law (p. 41), 
we obtain the expression: 

and, by the use of this, we can find the volume, v it that a given mass 
of gas will occupy at any desired temperature and pressure (7" 2 andp t ) 
from its volume, v t , at a given temperature and pressure (7\ and/?^. 

Example. A certain mass oj gas occupies 21 1 cm 3 at 18°C and 740 mm 

1 Directly proportional is the mathematical expression of the fact that the 
volume increases in the same proportion as the kelvin temperature increases. 


pressure. What volume will it occupy (still gaseous) at — 20°C and 
770 mm pressure? 

18°C = (18 + 273) K = 291 K 

-20°C = (-20 + 273) K = 253 K 

PlV i _ Pt» i 

T x " T, 

740 X 21 1 770 X v t 



v s = 

740 x 211 x 253 




At this stage, inspect your fraction to see if it agrees with what common 
sense would lead you to expect, 

Thus (a), the 21 1 of the numerator, is the original volume of the 
gas in cm 3 . The pressure is changing from 740 to 770 mm; that is, an 
increase of pressure. This should decrease the volume. The fraction 


— - is actually doing so and is therefore correct. 

(b) The temperature is falling from 291 K to 253 K; that is, the 

volume of the gas should be decreasing. The fraction — — decreases 

the volume as required and is, therefore, correct. 
Using logarithms, v t = 176 cm 3 

Standard temperature and pressure 

Since the volumes of gases change in such a marked manner with 
changes of temperature and pressure, it is necessary to choose a suit- 
able value of each as standards to which gas volumes can be referred. 
The standards chosen are 0°C and 760 mm pressure and these are 
known as standard temperature and pressure, usually contracted to 

That is, s.t.p. indicates standard temperature and pressure or 0°C 
(273 K) and 760 mm pressure (1 atm). In the past, the alternative, 
N.T.P. (normal temperature and pressure) was used, the values being 
the same. 

Example. A certain mass of gas occupies 146 cm 3 at li"Cand 738 mm 
pressure. Calculate its volume at s.t.p. 

18°C = 291 K s.t. = 273 K (0°C) 
s.p. ■» 760 mm 



Pl"l _ Pt«t 

T t " T t 

738 X 146 760 X v. 


v t = 


738 X 14 6 x 273 (Inspect this fraction 
760 > 

291 as described above) 

= 133 cm 3 

Dalton's Law of Partial Pressures 

This law states that, in a mixture of gases which do not act chem- 
ically together, each gas exerts a partial pressure which is the pressure 
it would exert if it alone filled the containing vessel at the same tem- 
perature and pressure. Then the total gas pressure is the sum of the 
partial pressures of the constituent gases. 

This law applies most commonly to the case of a gas collected over 
water. For example, suppose 100 cm 3 of an insoluble gas are col- 
lected over water at a barometric pressure of 745 mm and at 15°C. 
If it is saturated with water-vapour, the true pressure of the gas is 
(745—13) mm, since the vapour pressure of wate rat 15°C is 13 mm. 

732 273 
Dry, at s.t.p., the gas would occupy 100 x - x — - cm 1 . 

760 288 


Calculate the volumes which will be occupied at the given final tem- 
peratures and pressures by the gases whose initial volumes, temperatures 
and pressures are given. 


Initial T. & P. 

Final T. & P. 


273 cm' 

0"C and 760 mm 

14°C and 861 mm 


1638 cm' 

0°C and 819 mm 

!5°Cand 864 mm 

1000 cm' 


23°C and 800 mm 


500 cm' 


-48°C and 750 mm 


1000 cm' 




760 cm' 

27°C and 700 mm 


The exam 

pies below need the use 

of logarithms 


700 cm' 1 

17°C and 740 mm 


133 cm' : 

I4°C and 745 mm 

17°C and 750 mm 


55 cm' 1 

14°C and 744 mm 



574 cm' 1 


I5'C and 735 mm 


70 cm' ' 


18°C and 745 mm 


121 cm' 


120°C and 742 mm 


534 cm' 



Chapter 6 

Experimental Work on Formulae, 
Reactions and Equivalents 

THIS chapter deals with experimental work which can be used to 
determine some empirical (simplest) formula; of compounds and 
the numbers of various atoms involved in some displacement re- 
actions. Atomic weights of elements involved are assumed known 
and are quoted. 

Displacement of hydrogen by metal 

Many metals displace hydrogen from dilute acids. For experi- 
mental purposes, magnesium is convenient because it reacts rapidly 
with either dilute hydrochloric or dilute sulphuric acid at room tem- 
perature. The purpose of the present experiment is to determine 
(given Mg = 24 and H = 1 .008) what numbers of atoms of mag- 
nesium and hydrogen displace one another in the reaction. Mag- 
nesium is weighed normally on a chemical balance but hydrogen 
(being the lightest known gas) is more conveniently measured by 
volume and converted to weight from the fact that one dm 3 (litre) of 
hydrogen at s.t.p. weighs 0.09 g. 

Experiment. Weigh a small watch-glass, then put on to it about 
0.25 g of cleaned magnesium ribbon and weigh again. Obtain the 
weight of the metal by difference. Without the conical flask in posi- 
tion (Fig. 11), fill the siphon tube with water by blowing gently into 
the short tube attached to the aspirator. Close the clip and place the 
delivery tube in the measuring cylinder. (This procedure is necessary 
because at the end of the experiment the siphon tube will be filled 
with water and, if it were left empty at the start, the volume of water 
delivered would be too low.) Using a funnel (to keep acid away from 
the upper sides) pour about 30 cm 3 of dilute hydrochloric acid into 
a conical flask, carefully place the weighed magnesium as shown and 
connect the conical flask to the aspirator (supporting it suitably). 




Siphon Lube, 
full of water 



Fio. II. 
Displacement of hydrogen. 

Open the clip. If the apparatus 
is airtight, water will flow for a 
second or two, then stop. (If it 
continues longer, a leak exists and 
must be remedied.) Shake the mag- 
nesium into the acid. When all the 
metal has dissolved, allow time 
for the apparatus to cool. Then 
adjust the levels of water in the 
measuring cylinder and aspirator 
to equality (by raising or lowering 
one of the two as required). This 
gives atmospheric pressure in the 
aspirator. Read the volume of 
water in the measuring cylinder, after closing the clip and removing 
the cylinder. This volume is also the volume of hydrogen liberated. 
Room temperature and the barometer reading are required. 

Specimen readings 

Weight of magnesium — 0.260 g 
Volume of hydrogen = 252 cm 3 
Room temperature = 10°C 
Barometer reading = 752 mm (corrected for 

vapour pressure of water) 

1. Reduce the volume of hydrogen to s.t.p. 

From this, 

.Wi 752 X 252 760 X v t 

' T s 283 273 

752 x 252 x 273 _,. _ 
Vi ~- 283X760 =240 - 5cm3 

2. Find the weight in grams of this volume of hydrogen 
1000 cm 3 of hydrogen at s.t.p. weigh 0.09 g 
so 1 cm 3 of hydrogen at s.t.p. weighs 0.00009 g 

and 240.5 cm 3 of hydrogen at s.t.p. weigh 0.00009 X 240.5 g 

or 0.02165 g 
Approximating slightly on the experimental results: 0.0217 g of 
hydrogen are displaced by 0.260 g of magnesium. Assuming the 
atomic weights of hydrogen and magnesium to be 1.008 and 24 
respectively, the number of moles of hydrogen and magnesium atoms 
which replace one another are: 

and °™. or 0.0215 and 0.0.08 


24 * 



This is an almost exact ratio of 2 : 1 in moles, i.e., one atom of 
magnesium has displaced two atoms of hydrogen from the acid. This 
corresponds to the equation : 

Mg + H 2 S0 4 ->• MgSO« + H 2 

If desired, this experiment can be used to calculate the equivalent of 
magnesium as below. 
From the results: 
0.0217 g of hydrogen are equivalent to 0.26 g of magnesium 

1.008 g of hydrogen are equivalent to 0.26 X 


g of magnesium 


or 12.1 g. This (by definition, p. 37) is the gram-equivalent weight 
of the metal. 

Experiments on the formula of black copper oxide 

Black copper oxide contains only copper and oxygen. Evidence 
about its formula may be obtained either by analysing it by reduction 
in hydrogen (with heat) or by synthesising it by oxidation of the pure 
metal, both processes being quantitative. 

Method I. Reduction of the oxide 

Black copper oxide has been made by the chemical union of copper 
and oxygen. By the removal of oxygen in a stream of hydrogen the 
weights of copper and oxygen which were in combination can be 
determined (Fig. 12). 

(Coal-gas can be used instead of hydrogen for this experiment. Its 
use is not dangerous.) 

The tube must slope downwards towards A, otherwise water con- 
densed during the experiment might run back on to the heated part 
and crack the tube. 

Weigh a porcelain boat and weigh again with some pure dry copper 
oxide in it. Place the boat inside a hard glass tube. Generate hydrogen 
as in Fig. 12 and pass it through a calcium chloride tube to dry it. 
Allow the hydrogen to pass over the oxide until, when collected in a 
test-tube as shown, it burns quietly on exposure to a flame. This shows 
that all the oxygen in the apparatus has been expelled. Then light the 
hydrogen at the jet and warm the copper oxide. Soon a glow spreads 
through the oxide (an indication that chemical change is taking 
place). A reddish-brown powder, copper, is left, and drops of water 
collect at A. 
Allow the copper to cool in a current of hydrogen so that air 



cannot enter and so convert the red metallic copper into the oxide 

Weigh the boat and copper when cool. (The boat and contents 
should, to ensure complete reaction, be heated to constant weight.) 

Note that the copper oxide has been reduced by the hydrogen to 
red metallic copper, whereas the hydrogen has been oxidised to water. 
(See p. 162.) 

CuO + H, ->■ Cu -f H,0 

copper hydrogen copper water 

Cone. Hydrochloric 
- Acid 

Zinc and 

Calcium. Chloride 

to dry the 


Fig. 12. 
Reduction of copper oxide by heating in hydrogen. 

Specimen analysis 

Weight of boat = 4.32 g 

Weight of boat h oxide = 5.61 g 

Weight of boat + copper = 5.35 g 

.'. Weight of oxygen = 0.26 g 

.". Weight of copper = 1.03 g 

0.26 g of oxygen has combined with 1.03 g of copper 

Assuming the atomic weights of oxygen and copper to be 16 and 
63.5 respectively, the number of moles of oxygen atoms and copper 
atoms which combined together are: 

~\6 and 63~5 ° r °' 016 ( a PP rox -) in both cases 

This means that, in black copper oxide, copper and oxygen atoms 
combine together in equal numbers and the simplest formula of the 



oxide is CuO. This is copper(II) oxide in which the valency of the 

metal is 2. 

If desired, the equivalent of copper in this oxide can be calculated 

from the experimental results as below: 

0.26 g of oxygen are equivalent to 1.03 g of copper, 

1 03 X 8 
so 8 g of oxygen are equivalent to ' g of copper, or 31.7 g 


of copper. 

This (by definition, p. 37) is the gram-equivalent weight of copper. 
If coal-gas is used for this experiment, all details are the same 

except that the air can be swept out so quickly that there is no need 

to test for its removal. 

For class use, the following method is satisfactory: 

Blow a small hole in a test-tube and weigh the test-tube. Weigh 

again with some pure dry copper oxide in it. Push the rubber tube 

CuO + H, ->• Cu + H.O 




Dilute Nitric 

Fig. 13. 

Fig. 14. 

from the Bunsen burner into the test-tube, pass coal-gas through it 
and light it as shown in Fig. 13. 

Conduct the experiment in a similar manner to that described in 
the previous experiment and work out the result. 

Method 2. Oxidation method 

Weigh a small evaporating dish and clock-glass (7.5 cm diameter is 
suitable), and weigh again having added one or two small pieces of 
copper (not more than half a gram). Remove the clock-glass, add 
about 10 cm 3 of bench dilute nitric acid (approx. 4M), replace the 
clock-glass, and heat gently on a tripod and gauze in a fume-chamber 
(Fig. 14). There is a vigorous effervescence, brown fumes of nitrogen 
dioxide are seen, and the copper finally dissolves, giving a blue solu- 
tion of copper(ll) nitrate. 

3Cu + 8HN0 3 -* 3Cu(N0 3 ) 3 -f- 2NO + 4H 2 
2NO + O l -»-2NO, 



Continue to heat the solution (increasing the size of the flame to 
maintain a steady but not too vigorous evolution of vapour) until the 
whole of the excess nitric acid has been driven off and the copper(H) 
nitrate converted to black copper oxide. 

2Cu(N0 3 ) a -► 2CuO + 4NO s + O a 






Heat the dish very strongly for a few minutes to decompose any 
copper nitrate on the sides of the dish or on the clock-glass, allow to 
cool (in a desiccator if possible) and weigh the dish and clock-glass. 
Repeat the heating to constant weight. (The clock-glass minimises 
loss of liquid by spurting but does not materially reduce the rate of 

Specimen weighings 

Weight of evaporating dish and clock-glass = 20.210 g 
Weight of the above + copper = 20.634 g 

Weight of the above -f- copper oxide = 20.741 g 

Weight of copper used = 0.424 g 

Weight of oxygen combined = 0.107 g 

From these figures, 0.107 g of oxygen combined with 0.424 g of 

copper. Given that O = 16 and Cu = 63.5, 


moles of oxygen 

atoms combine with -j— - moles of copper atoms. These are equal 

numbers of atom moles at the value of 0.0067 approximately. That 
is, the simplest formula of the black oxide is CuO and it is copperiJI) 
oxide in which the valency of copper is 2. 

If desired, this experiment gives the equivalent of copper in this 
oxide, as below. From the results, 

0.107 g of oxygen are equivalent to 0.424 g of copper, 

so, 8 g of oxygen are equivalent to °' 4 q 4 * 8 -I of copper, or 

31.7 g. 

This (by definition, p. 37) is the gram-equivalent weight of the metal. 

Experiment on the reaction between zinc and 
copper(II) sulphate solution 

Crush a quantity of copper(H) sulphate crystals (about 5 g) in a 
mortar and dissolve them in water in a beaker, warming to hasten 
the solution. Weigh a small watch-glass, add a few strips of pure 

Plate 4. Welding 
chain links in the 
factory of Messrs 
Watson and Mac- 



zinc foil (weight about 1 g) and weigh again. The weight of the zinc 
can be obtained by difference. Put the zinc into the copper(ll) sulphate 
solution. Immediately the zinc becomes coated with a red film of 
copper and. on stirring, the copper falls to the bottom of the beaker. 
After a while the whole of the zinc will have disappeared and there 
will be a layer of red metallic copper on the bottom of the beaker. 
(More copper(II) sulphate crystals can be added if the colour indi- 
cates that the solution is dilute.) 

Filter off the copper and wash the small particles of copper adher- 
ing to the beaker into the filter-paper by means of a jet of water from 
a wash-bottle. Wash the copper several times with hot distilled water, 
and finally two or three times with methylated spirit (care being 
taken to extinguish any burners likely to set fire to the spirit). Allow 
the copper to dry and weigh it, together with the filter-paper, a clean 
filter-paper being placed on the right-hand pan of the balance and 
the weights on top of the filter-paper. 

Specimen calculation 

Weight of zinc used = 0.920 g 

Weight of copper displaced = 0.890 g 

Given Zn = 65 and Cu = 63.5, 


moles of zinc atoms displaced 

moles of copper atoms. That is, the number of moles of zinc 

and copper atoms concerned is equal at 0.014 approximately. This is 
in accordance with the equation: 

Zn + CuS0 4 -»- ZnSO« + Cu 

in which one atom of zinc replaces one atom of copper. (The reaction 
is really ionic as: 

Zn + Cu 2+ -> Zn a+ + Cu) 


1. Sketch and label an apparatus suitable for heating an oxide of 
copper in a current of dry hydrogen. State two precautions (with reasons 
for them) which are necessary for the safe conduct of the experiment. 
Describe what happens to the oxide of copper when heating begins. If 
0.429 g of oxide were reduced to 0.381 g of copper in an experiment, which 
oxide of copper (CuO or Cu s O) was used? (Cu = 63.5; O = 16.) 

2. Sketch an apparatus suitable for measuring the volume of hydrogen 
liberated by the action of a known weight of magnesium on excess of dilute 
hydrochloric acid. Indicate the precautions and readings necessary to 
obtain a reasonably accurate result. If0.200gof magnesium liberate 187 cm' 
of hydrogen (dry) at s.t.p., calculate the ratio of magnesium used to 

Right: The com 
plcted link 



hydrogen liberated in terms of gram-atoms. (Mg = 24; H = 1.008; 
1000 cm' of hydrogen at s.t.p. weigh 0.09 g.) 

3. A metal, X, precipitates copper quantitatively from copper sulphate 
solution. In an experiment, 0.480 g of X were found to produce 1 .270 g 
of copper after purification. If the equation for the reaction is: 

X + CuSO, — *- XSO, + Cu 

calculate the atomic weight of X. (Cu = 63.5.) 

4. 0.480 g of clean magnesium ribbon was slowly burnt in air with no 
appreciable loss of material. The cooled product was moistened with water 
(to destroy any magnesium nitride) and the resulting material was strongly 
heated to constant weight (0.800 g) of magnesium oxide. Given Mg = 24 
and O = 16, calculate the gram-atom ratio in which the two elements 
react and write the simplest formula of magnesium oxide. 

5. 1 .800 g of aluminium was dissolved with heat in slight excess of dilute 
hydrochloric acid. Aluminium hydroxide was precipitated by the addition 
of slight excess of ammonia solution and, after filtration, washing, drying 
and strongly heating to constant weight, 3.400 g of aluminium oxide 
remained. If the atomic weights of aluminium and oxygen are 27 and 16 
respectively, deduce from the above data the simplest formula of aluminium 
oxide. (Use gram-atom ratios.) 

6. When suitably heated in oxygen, lead can produce a red compound 
said to be Pb 3 0«. (This red compound can be entirely reduced to lead at 
red heat in a current of dry hydrogen.) Describe, in detail, an experiment 
by which (using this reduction) you would test the accuracy of the quoted 
formula. Outline any required calculation. (Pb = 207; O = 16.) 

7. Two oxides of lead, A and B, were heated in dry hydrogen to reduce 
them to metallic lead. In case A, 0.446 g of oxide left 0.414 g of lead; in 
case B, 0.717 g of oxide left 0.621 g of lead. Show that these figures are in 
accordance with the Law of Multiple Proportions and, given Pb = 207 
and O = 16, calculate the simplest formula: of A and B. 

8 0.108 g of a metal, M, of atomic weight 27, liberated 134 cm' of 
hydrogen (measured dry) from an acid, H a X, at s.t.p. Make the calculations 
(in gram-atom terms) which are needed to determine the simplest values 
for n and m in the equation: 

nM + twH.X — > M»Xm + mH, 
(H = 1.008; 1000 cm 3 of hydrogen at s.t.p. weigh 0.09 g.) 

9. 0.180 g of aluminium liberated, from dilute hydrochloric acid, 
235 cm 3 of hydrogen, measured (dry) at 10°C and 750 mm pressure. If the 
equation for this reaction is: 

Al + xHCl -»• Aid, 

+ l H * 

deduce (from the data given) the value of x. (Al = 27; H = 1.008; 
1000 cm 5 of hydrogen at s.t.p. weigh 0.09 g.) 

10 108 g of a metal, M, liberates 147 cm 8 of hydrogen from a dilute 
mineral acid, measured (over water) at 15°C and 745 mm pressure. Given 
also that hydrogen and oxygen combine in the volume ratio of 2 : 1 
respectively (at constant temperature and pressure), deduce by reference 
to gram-atom quantities the simplest formula of the oxide of M, assuming 
that the valencies of the elements concerned remain constant. (Atomic 



weight of M = 27; O = 16; 1000 cm 3 of oxygen weigh 1.44 g at s.t.p.; 
saturated vapour pressure of water at 15°C is 13 mm.) 

11. 0.360 g of metal, M, liberate from excess of dilute hydrochloric acid 
366 cm 3 of hydrogen (over water) at 15°C and 750 mm pressure. If, also, 
hydrogen and chlorine combine together in equal volumes at constant 
temperature and pressure and the valencies of all three elements remain 
constant, deduce (by reference to gram-atom quantities) the simplest 
formula of the chloride of M. (Atomic weight of M is 24; CI ■ 35.5; 
1000 cm 3 of chlorine weigh 3.20 g at s.t.p.; saturated vapour pressure of 
water is 13 mm at 15°C.) 

12. 0.934 g of a sample of granulated zinc containing a small percentage 
of non-reactive impurity yielded 356 cm 3 of hydrogen (over water) at 
740 mm pressure and 17°C by reaction with excess of dilute hydrochloric 
acid. By converting these quantities to gram-atoms, show that, with respect 
to zinc and hydrogen, these figures agree with the equation: 

Zn + 2HCI -+ ZnCl, + H, 
(allowing slight error for the impurity). Calculate the percentage of the 
impurity in the zinc (to one significant figure). (Zn = 65; H = 1.008; 
1000 cm 3 of hydrogen at s.t.p. weigh 0.09 g; saturated vapour pressure of 
water at 17°C is 14 mm.) 

13. A metal has an equivalent weight of 12. Calculate the weight of 
metal required to liberate from acid 525 cm' of hydrogen (dry) at 1 5°C 
and 750 mm pressure. (1 dm 3 or litre of hydrogen at s.t.p. weighs 0.09 g.) 

14. 0.162 g of a metal, X, were converted to 0.801 g of its chloride. 
Calculate the equivalent of X. If the valency of X is 3, what is its atomic 
weight? (CI = 35.5.) 

15. The oxide of a certain metal is completely decomposed into its 
elements by heat. Describe an experiment suitable to determine the volume 
of oxygen given off from a known weight of the oxide (mentioning pre- 
cautions needed for reasonable accuracy). If the volume of oxygen obtained 
from 1 . 72 g of the oxide is 94 cm 3 at 1 2°C and 750 mm pressure, what is the 
equivalent of the metal? (1 dm 3 or litre of oxygen at s.t.p. weighs 1 .44 g.) 

16. 2 g of a metal yielded 844 cm 3 of dry hydrogen from dilute acid at 
17°C and 765 mm pressure. Also, 4.20 g of the same metal were converted 
to 6.00 g of its oxide. Calculate the two equivalents of the metal, explain 
why they are different and use the results to illustrate the Law of Multiple 
Proportions. (Sec Q. 13 for hydrogen data.) 

Chapter 7 

Calculations Involving Weights 

WE have seen in Chapter 2 that every definite chemical reaction 
can be represented by means of an equation. In Chapter 4, the 
question of atomic weights was considered. In the present chapter 
we shall combine the knowledge obtained in both and show how a 
quantitative meaning can be assigned to an equation in terms of the 
commonly used weight units— more particularly the scientific unit, 
the gram. 

To calculate the molecular weight of a compound from its molecular 

We have noted that the molecular formula of a compound indi- 
cates the kind of atoms present in the molecule and their number; 
thus the formula H 2 S0 4 , for sulphuric acid, indicates that one mole- 
cule of the acid contains 2 atoms of hydrogen, 1 atom of sulphur and 
4 atoms of oxygen. The atomic weight of each of these elements is 
known and can be obtained from tables. The molecular weight of 
sulphuric acid can now be calculated by allowing the appropriate 
number of weight units for each element present and adding to 
obtain the total. (H = 1 ; S = 32; O = 16.) 
Thus H, S 4 

(2 X 1) + 32 + (4 X 16) 
= 2 + 32 + 64 
= 98 = molecular weight of sulphuric acid 

Taking another example: Calculate the molecular weight of red 
lead, PZ> 3 4 . (Pb = 207; O = 16.) 

Pb 3 4 

(3 X 207) + (4 x 16) 
= 621 + 64 
= 685 = molecular weight of red lead 

If the formula of the compound contains bracketed acid radicals, it 
will be simpler and more accurate for you to remove the brackets 




first and then proceed as above. Thus: Calculate the molecular weight 
of calcium nitrate, Ca(N0 3 ) 2 . (Ca = 40; N = 14; O = 16.) 
Ca(NO,) 2 
orCaN 2 0« 
40 + (2 x 14) + (6 x 16) 
= 40 + 28 + 96 

= 164 = molecular weight of calcium nitrate 
Later, with practice, you will be able to carry out the removal of the 
bracket mentally. 

It is also possible to calculate the weight of each element present in 
a given weight of compound from its formula. This information is 
usually stated as the percentage composition of the compound. 

To calculate the percentage by weight of each clement present in a 

compound from its formula 

Example: Calculate the percentage by weight of each element in 

calcium sulphate. CaSO A . (Ca = 40; S = 32; O = 16.) 

First calculate the molecular weight of calcium sulphate. 

CaS0 4 

40 + 32 + (4 X 16) 

= 40 + 32 + 64 

= 136. 

40 of these 136 units of weight are calcium, that is, the fractional 

weight of calcium in calcium sulphate is 40/136. Then the percentage 

weight is — x 100 or 29.4. 


Similarly the fractional weight of sulphur is — — and the percentage 


is ~ x 100 or 23.5. 

The percentage weight of the third element, oxygen, need not be 

calculated as it is given by the expression 100 — (% of calcium + % 

of sulphur), 

or 100 - (29.4 + 23.5) 

= 47.1 

These calculations can be set out compactly as below: Calculate 

the percentage composition of calcium hydroxide, Ca (OH) a . (Ca = 40; 

O = 16; H = 1.) 


or Ca0 2 H, 

40 + (2 x 16) + (2 X 1) 

= 40 + 32 + 2 

= 74 



40 40 

Calcium. Fractional weight — . Percentage weight — X 100 

74 74 

= 54.1 

32 32 

Fractional weight — . Percentage weight — X 100 
74 74 


i* /• 

= 43.2 
Hydrogen. Percentage weight = 100 — (54.1 + 43.2) = 2.7 
It is also possible to calculate the formula of a compound from its 
composition by weight. 

To calculate the simplest formula of a compound from its composition 
by weight 

This calculation is illustrated by the following worked example. 

Calculate the formula of a compound which has the following per- 
centage composition: sodium 43.4, carbon 11.3, oxygen 45.3. 
(Na = 23;C= 12; O = 16.) 

The fact that the atomic weight of sodium is 23 means that every 

23 parts by weight of sodium in the compound represent one atom of 


43 4 
Thus 43.4 parts by weight of sodium represent — ^- atoms of 

sodium. Similarly for the other elements present: 

11.3 parts by weight of carbon represent — '- atoms of carbon. 

45 3 
45.3 parts by weight of oxygen represent — '— atoms of oxygen. 


.*. Number of atoms represented is: 

sodium carbon oxygon 

43.4 11.3 45.3 

23 12 16 

or 1.89 0.94 2.83 

These cannot be the actual numbers of atoms present because 
fractions of atoms are impossible. We have to find the whole numbers 
which are in the ratio 1.89 : 0.94 : 2.83. To do this, divide all these 
figures by the lowest or, if this does not result in a whole number 
ratio, by the smallest difference. Then the number of atoms of each 
element is : 

sodium carbon oxygen 

L89 0.94 Z83 

0.94 0.94 0.94 

or 2 1 3 



That is, the formula is Na 2 C0 3 . 

The calculation is set out compactly in the following example: 

Calculate the formula of a compound which has the composition: 

magnesium 9.8%, sulphur 13%, oxygen 26%, water of crystallisation 

51.2%. (Mg - 24; S = 32; O = 16; H a O = 18.) 

% by weight 
Ratio of atoms or 

Divide by smallest (or 
smallest difference) 

Magnesium Sulphur 

g- 0-408 




g- 0.406 






26 - I fil 
16 = lM 





5 -!^ = 284 
18 *•** 




.'. the formula is MgS0 4 .7H s O 
Empirical and Molecular Formula; are discussed on p. 107. 


We have seen already that the equation: 

CuO + H,S0 4 -*■ CuSO« + 
(63.5 + 16) (2 + 32 + 64) (63.5 + 32 + 64) 
79.5 98 159.5 

H 2 

(2 + 16) 

means: 'One molecule of copper oxide reacts with one molecule of 
sulphuric acid producing one molecule of copper sulphate and one 
molecule of water.* 

The appropriate molecular weights having been inserted, as above, 
it also means that 79.5 parts by weight of copper oxide react with 98 
parts by weight of sulphuric acid, producing 159.5 parts by weight of 
copper sulphate and 18 parts by weight of water. These 'parts by 
weight' may be any desired weight-units grams, tons, ounces, 
pounds, kilograms— provided that the same unit is used throughout. 

Obviously, the figures given by the equation can be used to cal- 
culate any required information about the weights of the four sub- 
stances concerned. 

Example 1 . What weight of copper sulphate could be obtained by 
starting with 10 g of copper oxide? 

From the equation, 79.5 g copper oxide yield 159.5 g copper 

.". 10 g copper oxide yield 159.5 X -— g copper sulphate 
= 20. 1 g copper sulphate 





Example 2. What weight of pure sulphuric acid would be needed to 
react with 15 tons of copper oxide? 

From the equation, 79.5 tons of copper oxide need 98 tons of 
sulphuric acid. 

.". 15 tons of copper oxide need 98 X -— tons of sulphuric acid 

= 18.5 tons sulphuric acid 

This means that equations have now been given a quantitative 
meaning in terms of ordinary weight-units, instead of simply in terms 
of atoms and molecules. This makes them extraordinarily useful for 
making calculations of the weights of materials needed for chemical 
reactions and the weights of products obtainable. Chemical manu- 
facturers base all their calculations of weights of materials on 

We will now discuss the steps necessary in using an equation 
correctly for weight calculations of this type. 

A balanced equation is necessary 

It has already been explained (p. 23) that, by a balanced equation, 
we mean one which has the same number of each kind of atom on the 
right of the equation as on the left. An unbalanced equation implies 
that atoms have been created or destroyed; it is therefore wrong, 
and calculations based on it are certainly unreliable. The first and 
absolutely essential step, then, is to obtain a balanced equation. 

Facility in producing balanced equations is only attainable with 
practice, but this absolutely inviolable rule must be remembered: 
The formula of a compound is absolutely fixed and unalterable and an 
equation must be balanced by taking appropriate numbers of molecules 
of the substances concerned, not by attempting alteration of their 

The following brief account will illustrate the process of balancing. 

To obtain a balanced equation for the action of hydrogen sulphide, 
// 2 S, on sulphur dioxide, S0 2 , producing water, H s O, and sulphur, S. 

The skeleton, but unbalanced and incorrect, 'equation' will be: 

H 2 S + S0 2 -»- H 2 + S (UNBALANCED) 

We cannot alter any of these formulae or the symbols. 

Now, both the oxygen atoms of the S0 2 molecule form water, 
therefore 2H 2 must be obtained. This gives: 

H 2 S + SO a -»- 2H s O + S (UNBALANCED) 
The 2H a O on the right now requires 2H 2 , which must be provided 
by taking 2H 2 S. This gives: 

2H 2 S + S0 2 -»- 2H 2 Q + S (UNBALANCED) 

We now have, on the left, 2S from the 2H 2 S (remember the 2 
multiplies all the H 2 S) and S from the S0 2 , therefore we must have 
3S on the right and the balancing is complete. 

2H 2 S + S0 2 -»-2H 2 + 3S 

Some such balancing process is necessary for all equations, but 
many of them become so familiar with frequent use that they can be 
set down correctly at once. 

Insertion of molecular weights into the equation 

Here, it is very desirable to remember that it is unnecessary to 
insert the molecular weights of any materials unless they are actually 
concerned in the calculation you are performing. 

Consider, for example, this problem. 

Calculate the weight of calcium nitrate which would be formed by 
treating 148 grams of slaked lime, Ca(OH) 2 , with excess of dilute nitric 
acid. (Ca = 40; O = 16; H =1.) 

Here the balanced equation is: 

Ca(OH) 2 + 2HN0 3 ->- Ca(NO s ) 2 + 2H 2 

slaked nitric calcium water 

lime acid nitrate 

In the problem, the only two substances mentioned quantitatively 
are calcium nitrate, of which a weight is to be calculated, and slaked 
lime, of which the weight is given. (Nitric acid is only mentioned as 
'excess'.) The only molecular weights we need insert are, therefore, 
those of calcium nitrate and slaked lime. 2HN0 3 and 2H 2 may be 
ignored once the equation has been balanced, because the problem is 
not concerned with them. 

So we get, 

Ca(OH) 2 + 2HN0 3 -*- Ca(N0 3 ) 2 + 2H 2 
40 + 32 + 2 40 + 28 + 96 

74 164 

From the equation, 74 g slaked lime yield 164 g calcium nitrate. 
.'. 148 g slaked lime yield 164 X — g calcium nitrate 

= 328 g calcium nitrate 

Calculate the weight of lead which would be obtained by heating 
34.25 g of red lead in a stream of hydrogen and the weight of water 
formed at the same time. (Pb = 207; H = 1 ; O = 16.) 

Writing the balanced equation and inserting the molecular weights 
of the materials concerned in the calculation, we have : 


Pb 3 4 + 4H 2 -»- 3Pb + 4H,0 
red lead 

621 + 64 4(2 + 16) 

685 621 72 

(i) From the equation, 685 g red lead yield 621 g lead. 

.". 34.25 g red lead yield 621 X 


g lead 


= 31.05 g lead 
(ii) From the equation 685 g red lead yield 72 g water. 

.-. 34.25 g red lead yield 72 x 


g water 

= 3.6 g water 


All Atomic Weights required for these calculations can be found on 
p. 532. 

1. How many tonnes of copper could be obtained by displacing copper 
from copper sulphate solution by 16.25 tonnes of zinc? 

2. What weight of sodium oxide, Na s O, could be made from 1.15 g 

3. Find the empirical formula: of the following compounds from their 
compositions by weight: 

(«) Zn, 47.8%; CI, 52.2% 

(b) Na. 39.3%; CI, 60.7% 

(r) Cu, 39.5%; S, 20.3%; O, 40.2% 

(«/) Pb, 62.5%; N, 8.45%; O, 29.05% 

4. Calculate the percentage by weight of each element in the following 

(a) Sodium hydrogen carbonate, NaHCO, 

(b) Calcium chloride, CaCI, 

(c) Ammonium sulphate, (NH 4 ) 3 SO, 

(d) Sodium thiosulphate, Na 2 S 8 O a 

5. What weight of dilute nitric acid (containing 10% of the pure acid) 
will be required to dissolve 5 g of chalk, calcium carbonate? 

6. How many grams of hydrogen sulphide would be necessary to 
precipitate 7.5 g of copper sulphide from a copper sulphate solution? 

7. 76.5 g of sodium hydrogen carbonate were heated strongly. What 
weight of carbon dioxide was obtained ? If a dilute acid had been added, 
what weight of carbon dioxide would have been obtained in this case? 

8. What weight of nitrogen dioxide could be obtained by heating 11.1 e 
of lead nitrate? 

9. 50 g of ammonium chloride were heated with 40 g of calcium 
hydroxide. What weight of ammonia gas would be evolved ? Which of the 
reagents is in excess and by how much? 



10. How many ounces of anhydrous zinc sulphate (ZnSO,) would be 
formed on completion of the reaction between 2 oz of zinc and dilute 
sulphuric acid containing 2 oz of the pure acid (HjSOJ? (N.U.J.B.) 

1 1 . How many grams of hydrochloric acid, containing 20% by weight of 
hydrogen chloride, would be required to dissolve 13 g of zinc? (N.U.J.B.) 

12. In a determination of the equivalent weight of carbon, 0.74 g of 
the element was burnt in a current of oxygen, the products of combustion 
were passed over heated copper(II) oxide, and the resulting carbon dioxide 
was absorbed in potash bulbs. The increase in weight of the potash bulbs 
was 2.69 g. Calculate to two places of decimals the equivalent weight of 

What was the object of using copper(II) oxide in this experiment? (C.) 

13. A compound has the percentage composition N = 19.31%, 
Ca = 27.58%, CI - 48.96%, and H = 4.13%. Calculate (a) the simplest 
formula for the compound, (6) the volume which the nitrogen, present in 
14.5 g of the compound would occupy at s.t.p. 

Atomic weights: N = 14, Ca = 40, CI = 35.5, H = 1. A molecular 
weight in grams of a gas occupies 22.4 dm 3 at s.t.p. (D.) 

14. Give a brief account of the chemical reactions involved in the 
extraction of iron from its ores. 0.1867 g of a sample of iron containing 
carbon as an impurity was dissolved in dilute sulphuric acid, filtered, and 
the filtrate heated with a slight excess of concentrated nitric acid. An excess 
of ammonium hydroxide solution was then added to the solution and the 
resulting precipitate was filtered off, washed, dried, and finally heated to 
redness until the weight was constant. The weight of the product was 
0.2600 g. Give the reactions involved in this process and calculate the 
percentage of iron in the original sample. (Fe = 56.) (L.) 

15. Pure calcium carbonate contains 44% by weight of carbon dioxide. 
Some dried chalk weighing 0.4 g was dissolved in dilute hydrochloric acid. 
The carbon dioxide given off had a volume of 85.5 cm* measured at 
750 mm pressure and 12°C. Find: 

(o) The volume of the carbon dioxide at s.t.p. 

(b) The weight of the carbon dioxide. 

(c) The percentage of pure calcium carbonate in the chalk. 

(H = 1,C = 12,0 = 16. Weight of 1 dm' of hydrogen at s.t.p. = 0.09 g.) 

Chapter 8 

Atomic Structure; Shape of Molecules; 
Kinetic Theory; Periodicity 

IT seems fairly certain that for most of the nineteenth century, 
atoms were regarded as very small spherical particles like a very 
minute lead shot. It was believed that no smaller particle could 
exist, and that atoms were solid and homogeneous. This state of 
affairs has been very greatly changed in recent years, mainly by the 
pioneer work of Lord Rutherford. 

It is now believed that atoms are themselves built up from three 
smaller particles— the proton, the electron and the neutron. The 
proton is a positively charged particle of mass about equal to that of a 
hydrogen atom. The electron is negatively charged, its charge being 
equal but opposite to the charge on a proton. It has a very small 
mass, about 1/1850 of the mass of the proton. The neutron has no 
charge, and its mass is about equal to the mass of a proton. 

Charge Mass ("C = 12) 

proton -f- 1 1 

electron —1 1/1850 

neutron nil 1 

Discovery of these particles 

Electrons. In the late nineteenth century, a great deal of work was 
done on the effects of electrical discharge at very high voltage (by 
induction coil) through elementary gases at very low pressure. This 
led to the discovery of cathode rays. These emerge at right angles to 
the cathode and travel in straight lines. If passed through an electro- 
static field, they are deflected away from the negative plate, i.e., they 
are negatively charged. The rays can exert mechanical pressure and 
convey substantial amounts of kinetic energy so that a metallic 
object on which they impinge will be heated and may even dis- 
integrate. From these facts, it was concluded that cathode rays 
consist of a stream of negatively charged particles in rapid motion 
(about 10» cm/sec). These particles were given the name of electrons 




They are produced by all known gases and have been shown to have 
a mass of 9.1 X 10 -28 g and a charge of 1.6 X lO -10 coulombs. 

At the same time, positively charged particles leave the area of 
the anode, their nature depending on the identity of the gas used. 
Hydrogen, for example, yields hydrogen ions, H + . The effect of the 
discharge is to ionise hydrogen atoms into electrons and hydrogen 

Protons. Rutherford carried out experiments early this century in 
which hydrogen was bombarded by fast alpha-particles (helium ions, 
He 2+ ) from a radioactive source. He found that very penetrating 
particles were produced, of approximate mass 1 ( 12 C = 12) and carry- 
ing an electrical charge equal to that of the electron, but positive. The 
particles were named protons and are, in fact, the same as the posi- 
tive particles produced in a hydrogen discharge tube (above). Alpha- 
particles ionised hydrogen atoms by knocking out electrons from 

H-*H + -f-<?- 

Neutrons were discovered (1930) as a very penetrating radiation 
knocked out of boron nuclei when this element was subjected to the 
action of alpha-particles from the radioactive element, polonium. A 
neutron has very nearly the same mass as a proton but no electrical 
charge. This lack of electrical properties explains why neutrons were 
detected so much later than electrons and protons. 

Arrangement of these particles in the atom. Nuclear Theory 

In 1906, Rutherford noticed that if a stream of alpha-particles was 
passed, as a very thin pencil, through gold leaf (of thickness about 
one-millionth of a centimetre) and then on to a photographic plate, 
a certain scattering of the alpha-particles was evident. Later (1909) 

K- particles 

Gold leaf 



Fio. 15. 

Effect of gold leaf on alpha-particles. 



Geiger and Marsden studied this scattering more accurately and 
extended the observations by using a zinc sulphide screen on a 
rotating arm. The screen scintillates when an alpha-particle strikes it. 
It was found that the straight-line path of the great majority of the 
alpha-particles was little affected by the gold leaf but about one in 
8000 of them was deflected through an angle of 90° or more from its 
original direction. Rutherford commented later that, considering the 
thinness of metal used and the high velocity and mass of the alpha- 
particles, these large deflections were 'about as credible as if you had 
fired a fifteen-inch shell at a piece of tissue paper and it came back 
and hit you'. Rutherford deduced that these few large deflections of 
alpha-particles required atoms of gold to contain a region (the 
nucleus) which: 

1. is positively charged because it repels the positively charged 
alpha-particles ; 

2. is relatively massive and highly charged because the deflections 
are large; 

3. must occupy a very small space because so very few alpha- 
particles are materially affected. 

According to the current version of this nuclear theory of Ruther- 
ford, protons and neutrons of an atom are concentrated into the 
nucleus which is, therefore, positively charged, relatively massive but 
minute. The electrons, equal in number to the protons, and so making 
the atom electrically neutral, are outside the nucleus. The whole bulk 
of the atom, defined by the outermost electron ring, is very great 
compared with that of the nucleus -in Rutherford's analogy about 
the same proportion as the dome of St Paul's Cathedral to a man's 
clenched fist. 

Arrangement of electrons in the atom 

Bohr (1913) put forward a theory of electron positioning which is 
still generally accepted for chemical purposes. It was developed 
originally in connection with the hydrogen atom, which contains only 
one proton (as nucleus) and one electron. Bohr suggested the exist- 
ence of certain circular orbits (or shells) at definite distances from the 
nucleus in which the electron may rotate. These orbits are associated 
with definite energy content of the electron, increasing outwards 
from the nucleus. While the electron adheres to any one orbit 
(usually the innermost, of lowest energy), the atom radiates no 
energy. If, however, the atom absorbs energy, the electron may jump 
to one of the outer orbits of greater energy. If it later falls back to 
the inner orbit of lower energy, energy will be radiated as light of a 
definite colour (or frequency). Thus, if several outer, higher energy 
orbits are involved, an optical spectrum should be obtained, showing 



several lines at definite frequencies. Bohr was able to calculate the 
theoretical frequencies for such a spectrum and show that they 
accorded well with observed hydrogen spectra. It is also known that 
spectra of other elements indicate similar electronic orbits in their 

It was known that many cases occurred in which pairs or triplets 
of spectral lines occurred close together, e.g., the two yellow lines of 
sodium. To explain these slightly varying energy levels in the same 
electron shell, it was suggested that some orbits are circular and 
some elliptical, but this feature can usually be ignored for chemical 
purposes, at any rate at the present level. 

This, and much subsequent work, leads to the following con- 

1. Several groups of electrons may occur in an atom and each 
group is known as an electron shell. Shells are numbered 1, 2, 3, etc. 
outwards from the nucleus. All electrons in a given shell have 
approximately equal energy. This energy increases in successive 
shells outwards from the nucleus. 

2. The maximum possible number of electrons in a shell numbered 
n is 2n 2 , i.e., in successive shells 2, 8, 18, 32, . . . electrons. 

3. In the outermost shell of any atom, the maximum number of 
electrons possible is 8. 

A diagrammatic representation of a typical atom would be as 
shown in Fig. 16. 


\\\ Protons 
1 12 Neutrons 

electrons revolving in orbits 2.8.1 

Fig. 16. 

Sodium atom (diagrammatic). 


The simplest atom is that of hydrogen. It consists of a single 
proton with one electron rotating round it. In order of complexity, 
the simpler atoms are made up as in the following table, neutrons 
being omitted for reasons explained later: 
























































































8 1 





8 2 

The number of protons on the nucleus (which equals the number of 
electrons in the shells) is called the Atomic Number of the element 

Types of chemical combination 

From the table above, it will be seen that neon and argon have 
eight electrons in their outermost electron layer, i.e., an octet of 
electrons. This structure is very stable and extremely difficult to 
disturb. In consequence, these two gases are chemically inert and 
form no compounds with other elements. They are self-satisfied. In 
the simpler rare-gas, helium, the duplet of electrons is equally stable 
and functions like the octet. 

The tendency of other elements is to try to attain this rare-gas 
structure of a stable outer octet (or duplet) of electrons and their 
chemical behaviour is a reflection of this tendency. On this general 
principle, elements combine in two main forms of combination, 
known as the electrovalent (or polar-valcnt) and covalent types. 

Elcctrovalent Combination 

(Also known as polar-valent.) In this type, a metallic element or 
group loses, from its outermost electron shell, a number of electrons 
equal to its valency. These electrons pass over to the outer electron 



shells of non-metallic atoms with which the metal is combining. By 
this means, an electron octet is left behind in the metal and created in 
the non-metal. Both elements now have the outer electron structure 
of a rare-gas, but the metallic particles have a positive charge from 
the excess proion(s) left on the nucleus, while the non-metallic 
particles are negatively charged from the added electron(s). The 
particles are then known as ions. 

Before combination 

After combination 

Sodium atom 
Protons Electrons 
11 2,8,1 

11 2,8 

Sodium ion -f- 

Chlorinc atom 

Protons Electrons 

17 2,8,7 

17 2,8,8 

Chlorine ion — 

Both ions now possess stable outer electron octets, like a rare-gas. 

No molecules of sodium chloride are formed. Because of the 
attraction of the oppositely charged Na + and Cl~ ions for one 
another, the ions arrange themselves into a rigid, solid shape called a 
crystal, but they remain quite separate. No molecules of sodium 
chloride form. The combination can only be expressed in ionic form 
as Na + Cl~, meaning an association of sodium and chlorine ions in 
equal numbers. For the sodium chloride crystal, see p. 371. 

Characteristic properties of electro- (or polar-) valent compounds 

(1) Polar compounds do not contain molecules. They consist of 
aggregates of oppositely charged ions. In consequence, if they are 
melted, or dissolved in water, to make the ions mobile, they conduct 
electricity and are, therefore, electrolytes (Chapter 13). 

(2) They are solids and do not vaporise easily. 

(3) They will not usually dissolve in organic solvents such as 
toluene, ether, benzene, etc. 

Salts, alkalis and bases arc electrovalent and acids, when in solu- 
tion in water, also show electrovalency. 

The following cases are given in further illustration of electro- 

Calcium chloride 

In the calcium ion, the two excess nuclear protons produce a 
double positive charge; in each chlorine ion, the excess electron pro- 
duces a single negative charge, i.e., Ca 2+ .2C1 _ . 


Before combination 

After combination 

Calcium oxide 

Before combination 

Calcium atom 

Protons Electrons 
20 2,8,8,2 

Two chlorine atoms 

Protons Electrons 
17 2,8,7 

17 2,8,7 

These valency electrons pass to the 
chlorine atoms 

Calcium ion -f -f- 

Protons Electrons 
20 2,8,8 

Calcium atom 

Protons Electrons 
20 2,8,8,2 


Two chlorine ions - 

Protons Electrons 
17 2,8,8 

17 2,8,8 

Oxygen atom 

Protons Electrons 
8 2,6 


These valency electrons pass to the 
oxygen atom 

After combination 

Calcium ion -f- + 

Protons Electrons 
20 2,8,8 

Oxygen ion 

Protons Electrons 
8 2,8 

As stated on p. 67, the calcium ion acquires a double positive 
charge; the two excess electrons produce a double negative charge 
on the oxygen ion, i.e., Ca 4+ .O s_ . 

Note the presence of the outer octet of electrons in all the above 

Covalent Combination 

In this type of valency, electrons are not actually gained or lost by 
the atoms concerned. They pass into a 'shared' state. 

Consider two chlorine atoms. Each has the electron structure 2,8,7. 
In covalency, the atoms contribute one electron each to a 'shared- 
pair'. In this way, both obtain an approximation to the external 
octet by making fourteen electrons do the work of sixteen (Fig. 17). 

Here, actual molecules are produced, not ions. Each 'shared-pair' 


electron passes from an orbit controlled by the nucleus of one 
chlorine atom into an orbit controlled by the nuclei of both chlorine 
atoms. This joint control of the orbits constitutes the valency bond. 
Other examples of molecules formed by covalency are shown in 
Fig. 1 8. Each shared electron pair is made up of one electron from 
each atom concerned. All the atoms obtain a shared octet or duplet 
of electrons. Each shared pair is represented by a stroke in the con- 
ventional formula, which is given for ethanol in illustration. 

Two separata Chlorine atoms:- 

(M/e/eus(ft)/7 protons 
i and 18 (or 20) neutrons 
(£/ectroia(o) 2-8-7 

Chlorine molecule CI-CI 


FlO. 17. 

Chlorine atoms and molecule (diagrammatic). 

Characteristic properties of covalent compounds 

(1) Covalent compounds consist of molecules. They contain no 
ions, are unable to conduct electricity and so are non-electrolytes. 

(2) Simple covalent compounds are gases or volatile liquids, e.g., 
ammonia, carbon dioxide, ethanol (alcohol). This is so because their 
molecules are electrically neutral and have little attractive force for 
each other. In more complex covalent molecules, e.g., naphthalene, 
the atomic nuclei (+) of one molecule and the electrons (— ) of 
another attract each other. As the molecules come together, the 
electrons of each begin to exert repulsive forces on each other. The 
forces of attraction and repulsion are balanced in the formation of a 
crystal. These van der Waals forces are, however, rather weak and the 
crystals have low melting-points (e.g., naphthalene 81°C) compared 
with ionic crystals (e.g. Na+Cl", 801°C). 

(3) Covalent compounds are usually soluble in covalent organic 
solvents, such as benzene or carbon disulphide. 

Another variety of covalent bond (i.e., electron sharing) has been 
given the name of co-ordinate bond. This bond is characterised by the 
fact that the two shared electrons are both supplied by one of the 


-H Jt H 











H— C— C— O-H 

H H 

• electron of H 
o electron of othe 






Fig. 18. 

Some covalent molecules. 

participating atoms (not one electron by each atom as in an ordinary 
covalent bond). A co-ordinate bond is formed (in simpler cases, at 
least) when one of the participants possesses a lone pair of electrons, 
i.e., a pair not directly concerned in its existing valency bonds. This 
lone pair is donated to an atom needing them to build up, or com- 
plete, an electron octet or duplet of great stability. The ammonia 
molecule (Fig. 19) possesses such a lone pair of electrons; it can exer- 
cise co-ordinate valency towards a hydrogen ion (proton) from an 
acid to produce the ammonium ion, NH 4 + . This bonding supplies an 

o electron of N 

• electron of H 


r h 


r h ~i 


H? N J + H+-*- 

HI N 5H 


H — N-MH 





ammonia proton ammonium ion 

Fio. 19. 


immonium ion. 


electron duplet to the hydrogen nucleus while still maintaining 
(though with sharing) the electron octet of the nitrogen atom. The 
proton, combining with the NH 3 molecule, carries over its positive 
charge to give the ammonium ion, NH 4 + . The co-ordinate bond 
is often indicated by an arrow pointing from donor atom to 
acceptor atom as shown, though there is certainly some internal 
equalisation of the four bonds of NH« + . This ion is associated 
(electrovalently) with some anion as a salt, e.g., NH 4 + C1 _ or 
NH 4 + NO,-. 

Again, the ammonia molecule participates by co-ordinate bonding 
in the formation of the teirammine copper(Il) ion (also called 











H— N 

1 t 




N— H 


tetramminecopper (II) ion 

o electron of N 
• electron of H 

Fio. 20. 
Tctramminc copper(lI) ion. 

cuprammonium ion) which gives the intense blue coloration when 
excess of ammonia is added to copper sulphate solution. In this 
copper(II) sulphate, copper is in the 2,8,17,2 electron state with the 
outermost 2 (valency) electrons transferred to the SO* 2- ion and the 
outer electron shell vacant in the ion, Cu* + . Four NH 3 molecules 
form co-ordinate bonds with the Cu i+ ion by means of their lone 
pair electrons to create a shared electron octet in the vacant valency 
electron shell of the copper(II) ion (Fig. 20). 

Oxygen, having a 6-electron outer shell and accepting a lone pair 
to complete the shared octet, is a frequent participant in co-ordinate 
bonds (as acceptor atom). A simple example of this is seen in the 
relation of phosphorus trichloride to phosphorus oxychloride. 







(3 + (°)- a Qys° ra - 






o electron of P 

• electron of Cl or 

Fig. 21. 

Phosphorus oxychloride. 

Compounds containing only co-ordinate bonds and covalent bonds 
are very similar in properties to purely covalent compounds. Both 
types are non-electrolytes (having no ions); the simpler examples are 
usually liquids in ordinary room conditions, but, with co-ordinate 
bonds present, tend to be less volatile. 

Metallic bond 

Metals are held together in solid crystalline form by metallic bond- 
ing in something like the following way. The outer (valency) electrons 
of each atom are only loosely held, being relatively distant from the 
nucleus, and they separate from particular nuclei to move at random 
through the crystal lattice. The residual ions, now positively charged 
by loss of valency electrons, tend to repel each other but are held 
together by the moving electron cloud and some overlapping of 
residual electron orbits. 

This type of bonding is very strong in some metals, e.g., iron, 
which are difficult to shatter, but it is much weaker in, e.g., sodium 
or potassium, which can be cut with a knife. If an electric potential 

+ atomic nucleus 

Fio. 22. 

Metallic bonding. 



difference is applied to the ends of a metallic rod, the free electrons 
lose their random motion and move towards the positive end of the 
rod, being steadily replaced by more from the source of P.D. Hence 
metals are good electrical conductors. As freely moving electrons can 
convey heat energy, metals are also good conductors of heat. 

Since the bonding agent in a metal is mainly a moving electron 
cloud, the ions of most metals will usually slide relative to one 
another, under stress, without shattering the lattice and produce a 
new position of stability. This accounts for the malleability and 
ductility of many metals, though special temperature conditions 
may sometimes be necessary. 

Shapes of some Molecules 
Carbon dioxide 

In a carbon dioxide molecule, the carbon atom is bonded to each 
oxygen atom by two pairs of electrons in covalency. These electron 
groups exercise mutual repulsion and produce a linear molecule. 


o o 


o S C 



2 o 

• 00 

o electron of 
• electron of C 

Linear molecule 
of C0 2 

Molecular layer 
in cubic crystal 

Fio. 23. 
Carbon dioxide molecule and crystal. 

In solid carbon dioxide, 'dry ice', these linear molecules take up a 
cubic formation (bonded by van der Waals forces). One layer of 
molecules in the cube is shown. 



In a molecule of steam, the two hydrogen atoms arc bonded to the 
oxygen atom by shared electron pairs and the oxygen atom has two 
lone pairs of electrons not concerned in valency bonding. The two 
lone pairs exercise a stronger mutual repulsion than that between 


o electron of o 
• electron of H 


Bond angle 

Fig. 24. 
Molecule of steam. 

Molecule of 

any other electron pairs present. The resultant effect of this is to 
produce a 'bent' structure in which atomic nuclei all lie in the same 
plane with 104±° as the H-O-H angle. If the four electron pair 
repulsions were all equal, this angle would be 109±°. 


An ammonia molecule contains three hydrogen atoms bonded 
covalently, i.e., by shared electron pairs, to a nitrogen atom, which 

h h 

o electron of N 

• electron of H 





Bond angle in triangular pyramidal' 

Fig. 25. 
Ammonia molecule. 

has also one lone pair of electrons. The mutual repulsion of the four 
electron pairs produces a 'triangular pyramidal' shape of molecule 
in which the H-N-H angle is 107° (not 109$° which would be pro- 
duced if all four electron pair repulsions were equal). 




Iodine produces a molecule, I 2 , by formation of a covalent (shared 
pair) bond to which each iodine atom contributes one electron (so 
completing, for each atom, a shared electron octet). By the operation 
of van der Waals forces, the molecules then form crystals in which 


Covalent bonding in I2 molecule 
{Valency electrons only) 

Shape of 

Layer of I 2 molecules in crystal. 
Next layer covers dotted lines 

Fio. 26. 
Iodine molecule and crystal. 

they lie in a herring-bone pattern which is reversed in successive 
layers. One layer of iodine molecules is shown in the square. In the 
next layer, molecules lie over the dotted lines. 


The four covalency bonds of the carbon atom in methane are dis- 
tributed symmetrically in three dimensions, the angle between any 
pair of valency directions being 1094°. Each bond represents a shared 
electron pair to which a carbon and hydrogen atom supply one 
electron each. The molecule of carbon tetrachloride is similar with 
substitution of four chlorine atoms for the four hydrogen atoms of 






o electron of C 

o electron of H 



Shape of 

bond directions 

FlO. 27. 
Methane molecule. 

In the ethylene molecule, C 2 H 4 , the two carbon atoms are com- 
bined by two covalcnt bonds, i.e., two shared electron pairs. This 
indicates unsaturation in the molecule. Each carbon atom is also 







o electron of C 
• electron of H 

structural formula 
Fio. 28. 
Ethylene molecule. 

Shape of 

combined covalently to two hydrogen atoms. The nuclei of all six 
atoms involved lie in the same plane, i.e., the molecule is planar. The 
H-C-H angles are 117°. 


In the acetylene molecule, C 2 H a , the two carbon atoms are 
covalently combined together by three shared electron pairs. This 


H^cic J H 



o electron of C 
• electron of H 

H— C=C— H 

structural formula 

Fio. 29. 
Acetylene molecule. 




corresponds to a high degree of unsaturation in the molecule. Each 
carbon atom is also combined covalently with one hydrogen atom. 
The molecule is linear. 


The above discussion of valency and chemical combination con- 
tains no mention of neutrons. They appear to have little influence on 
the chemical properties of the atom. Their chief function is to help to 
determine the weight of the atom. For example, the sodium atom, 
with 1 1 protons and 12 neutrons, has a total weight of 23 units. The 
1 1 electrons are relatively negligible in weight. 

Many cases occur, however, in which two atoms contain the same 
number of protons but differing numbers of neutrons. Having equal 
numbers of protons, these atoms must also have equal numbers of 
electrons. These are arranged in the same way and give the atoms 
identical chemical properties. But the differing numbers of neutrons 
cause the atoms to have different weights. An element showing these 
characteristic properties — that is, possessing atoms of similar chemi- 
cal properties but different weights — is said to show isotopy and the 
varieties of the atom are called isotopes of the element. 

A well-known example of isotopy occurs in chlorine. 

Isotope 1 


17 protons 

18 neutrons 

35 total weight 


weights different 
CI = 35 


Isotope 2 

Nucleus Electrons 

17 protons .2,8,7 

20 neutrons 

37 total weight 

chemical properties identical 
CI = 37 

Dalton believed that all atoms of the same element were exactly alike. 
Isotopy has proved him wrong. But, except in a few exceptional cases, 
elements contain isotopes in almost constant proportions and so 
appear to act as if all their atoms are equal in weight. Chlorine, for 
example, is found to have its isotopes mixed in such proportions that 
its average chemical atomic weight appears constant. The lighter 
isotope predominates, giving a value of 35.5. 

It may also be mentioned that, in spite of recent, very prominent 
activities in splitting the atom, it can still be regarded as an in- 
divisible unit in chemical actions. Apart from modifications in the 



electron shells, an atom is conveyed as a whole unit in chemical 

Uranium has two principal isotopes. Both possess 92 protons and 
92 electrons. One isotope has 146 neutrons, giving U = 238; the 
other has 143 neutrons, giving U = 235. The lighter isotope con- 
stitutes about 0.7% of natural uranium. If the atom U = 235 acquires 
one extra neutron on the nucleus, it becomes unstable and divides 
into two approximately equal parts. This 'fission' is accompanied by 
a small diminution of mass, the combined final products having a 
mass slightly smaller than that of the uranium. Consequently, there 
is a very great evolution of energy, principally as heat, according to 
the Einstein equation, E = mc*, where E is the energy, m the mass 
lost and c the velocity of light. At the same time, neutrons are 
emitted. Other atoms of uranium absorb them and undergo fission, 
so that, if the mass of uranium is large enough, a chain reaction is set 
up, causing an atomic explosion. 

We have noted already that the number of protons in the nucleus 
of an atom is called the Atomic Number (Z) of the element concerned ; 
the sum of the protons and neutrons in the nucleus is called the 
Mass Number (A) of the element and it also expresses the mass of the 
atom (on the scale of '|C = 12) to the nearest integer. If an element 
is isotopic, it has as many mass numbers as isotopes. The atomic 
weight of an element for practical purposes is the weighted average 
of the accurate masses of its isotopes (on the same scale as above) as 
they occur in the element in experimental practice. For a given ele- 
ment, atomic weight is always a constant, or very nearly so, except 
for products of radioactivity, such as lead (p. 492). In some cases, 
recent atomic weight tables show variations of atomic weight caused 
by differences of isotopic distribution but, for ordinary purposes 
they are very slight, e.g., S = 32.064 ± 0.003. 




(Z) protons 

Mass Number 
(A) protons 
+ neutrons 

A -Z 


(2 isotopes) 


J|C1 35 
J?CI 37 


35.46(3 atoms 
37 approx.) 

If an element X does not show isotopy, in the symbol, £X, a is the 
Mass Number and b the Atomic Number of any atom of X. For a set 
of isotopes of X, b is constant and a varies to express the variable 
number of neutrons per atom. 



It was mentioned briefly at the beginning of Chapter 5 that a gas 
consists of molecules in rapid motion. The Kinetic theory makes the 
following assumptions about a 'perfect' gas. 

1. Molecules of a gas move in straight lines at very great velocity 
until they collide with each other or with the wall of the containing 
vessel. Gas pressure is exerted in this vessel as the result of collisions 
between the gas molecules and the containing wall. The number of 
collisions in unit time being very great, the pressure appears constant 
(at constant temperature). 

2. The total volume of the actual gaseous molecules is negligible 
relative to the capacity of the container. 

3. Forces of attraction or repulsion between the molecules of the 
gas are negligible. 

4. The average kinetic energy of the gas molecules measures the 
temperature of the gas. 

These assumptions are approximately fulfilled by real gases at 
ordinary temperature and pressure. 

Explanation of gas laws by the Kinetic Theory 

Boyle's Law. If the volume of a given sample of gas is reduced at 
constant temperature, the average velocity of the gas molecules remains 
constant so they collide more frequently with the walls of the smaller 
containing vessel. The more frequent collisions cause higher pressure. 

Charles's Law. A fall of temperature represents a decrease in the 
average kinetic energy of the gas molecules; that is average molecular 
velocity decreases (mass remaining constant). At constant pressure, 
this decreased velocity causes the sample of gas to occupy a smaller 

Law of Partial Pressures. In a non-reactive mixture, each gas 
exerts a separate pressure on the container because of collisions of 
its molecules with the containing walls. The total pressure on the 
container is caused by the sum of all the collisions. 

If gas pressure rises to very high values, molecules are crowded 
very much more closely together so that ultimately forces between 
them are no longer negligible and their volume becomes significant 
relative to the space occupied by the gas. Then the simple gas laws no 
longer apply. 


Substances in the solid state are made up of particles (atoms, mole- 
cules or ions) which are packed so tightly together as to leave only 



slight unfilled space in the mass of solid. (Contrast the very wide 
separation of gas molecules mentioned above.) The particles are 
bound together by forces strong enough to prevent movement of 
translation so that a solid has a definite shape which remains fixed 
(in constant conditions) unless sufficient force is supplied to shatter or 
distort the mass. Particles in a solid may however show a certain 
amount of movement of vibration. 

Particles involved in the assembly of solid structures may be the 

1. Atoms. Non-metals may form crystals by exercising covalency 
to combine together very large numbers of their own atoms. An 
example of such a solid is carbon as diamond, forming what is known 
as a macromolecule. For a full account of this, and of the spatial 
extension of the covalency of carbon that makes it possible, see p. 288. 

Metals form solid masses in a different way. Their atoms can quite 
readily part with valency electrons from their outermost shells, as in 
the following examples. 

Na-»Na + + e-; Cu -> Cu* + + 2e~ 

The resulting positively charged ions can then produce an ordered 
solid arrangement which depends on an equilibrium between their 
repulsive forces on each other and the binding efTect of the electron 
cloud which moves, continually and at random, among them. See 
also metallic bonding (p. 72). 

2. Molecules. Many cases occur in which non-metallic elements 
form molecules by covalent combination, e.g., iodine as I 4 . (For the 
similar formation of a molecule Cl„ see p. 69.) In appropriate condi- 
tions of temperature and pressure, the atomic nuclei (+) of one 
molecule and the electrons (-) of another molecule attract each other 
sufficiently to bring about a close approach. As the molecules come 
together, the electrons of each begin to exert repulsive forces on each 
other. The forces of attraction and repulsion are balanced in the 
formation of a crystal. These van der Waals forces are, however 
rather weak and the crystals tend to have low melting-points. Simple 
molecules of compounds, e.g., naphthalene, form crystalline solids 
by the operation of van der Waals forces in a similar way. For further 
account of the iodine molecule see p. 75. 

3. Ions. When elements combine by exercising electro-valency the 
resulting positively and negatively charged ions exercise powerful 
attractive forces on each other and can form a solid crystalline 
lattice. The arrangement of ions in such a lattice can usually be 
elucidated by X-ray diffraction. A well-known example is the lattice 
of solid sodium chloride formed from the ions, Na + and CI- for 
this see p. 371. Imperfect lattices sometimes occur. For example in 



ferrous sulphide, some of the iron positions may be unoccupied and 
this will cause the compound to deviate slightly from the Law of 
Constant Composition. 

Melting. If solids are heated, their constituent particles, atoms, 
molecules or ions, acquire greater kinetic energy and vibrate more 
violently. Eventually a point may come at which vibration overcomes 
the binding forces and the particles become mobile. The crystalline 
structure then collapses and a liquid state is reached in which the 
particles are free to move. The temperature, t"C, at which this occurs 
is called the melting-point of the solid and, at this temperature (and 
the prevailing pressure), the solid and liquid are in equilibrium. 
(Melting-points are generally only slightly modified, i.e., lowered, by 
moderate increase of pressure.) The energy which must be supplied 
to convert the solid at r°C to liquid at f°C is known as latent heat of 
fusion. Quantitatively, for example, 334 joules (80 cal) of heat energy 
convert one gram of ice at 0°C to one gram of water at 0°C (at 
standard pressure). If fairly simple molecules are only weakly bound 
together in crystals by van der Waals forces, the melting-point of 
the crystals will be quite low, e.g., naphthalene, C l0 H g , 8 PC. In 
some cases, e.g., iodine, some molecules may break away from the 
crystal directly into the vapour phase. This is known as sublimation. 
If binding forces are strong, e.g., in electrovalcnt lattices such as that 
of sodium chloride, or in some metallic crystals, melting-points will 
be much higher, e.g., for sodium chloride, 801°C, and for copper, 
1080°C, though sodium metal melts at only 97°C. When diamond is 
heated in vacuo (to prevent burning), it does not melt because the 
forces of covalency between the atoms are too great to be overcome 
by vibration with any ordinary source of heat. At very high tempera- 
ture, some carbon atoms will break away but by sublimation into the 
gaseous state, not by melting. (For determination of melting-point 
see p. 217.) 


It has already been mentioned that, when the particles which make 
up a solid acquire sufficient energy by heating, they become mobile, 
i.e., the solid melts and forms a liquid. The particles remain close 
together by virtue of the attractive forces between them but are 
perpetually in random motion. A liquid spreads to fill a containing 
vessel as far as possible but leaves a level liquid surface if the capacity 
of the vessel exceeds the volume of liquid. This occurs because the 
great majority of particles (usually molecules) at the surface are held 
by attraction from the mass of particles below them; only relatively 






few particles of liquid (those with more than average energy) can 

Because of the perpetual random motion of particles, usually mole- 
cules, in a liquid, very small particles of insoluble solids are kept in 
similar random motion in liquids because they are continually and 
unevenly bombarded by molecules of the liquid. This motion is called 
the Brownian Movement from its discovery by the botanist, Brown, 
in 1827, in connection with pollen particles in water. 

An ionic liquid is obtained when a crystal lattice of an electrovalent 
solid breaks down by the action of heat but the temperature required 
is usually far above room temperature, e.g., for K + C1 - , 768°C. An 
ionic liquid is an electrolyte, i.e., it conducts electric current by flow 
of ions to cathode and anode and is decomposed in the process, e.g., 

K+ + e~ -»■ K (at cathode); CI" -*■ $CI a + e~ (at anode) 

A molecular liquid contains covalent molecules, e.g., carbon 
tetrachloride, CC1 4 , chloroform, CHC1 3 , carbon disulphide, CS 2 . A 
great number of the simpler covalent compounds are liquid at room 
temperature and pressure, i.e., the melting-points of the correspond- 
ing solids are below room temperature. Such liquids are non- 
electrolytes; they do not conduct electricity having no content of 

Vapour pressure. All molecules in a liquid are subject to attractive 
forces from neighbouring molecules. For molecules in the body of 
the liquid, these forces balance each other but, for molecules at the 
surface, there is a resultant attractive force acting downwards. In 
spite of this, some molecules of greater than average energy can 
escape from the surface of the liquid into the surrounding space and 
are lost. The liquid is said to evaporate. Escape of these particles of 
high energy lowers the average kinetic energy of those remaining; 
consequently, the temperature of the remaining liquid falls. This 
coldness can be felt if a small quantity of a liquid such as ether is 
allowed to evaporate on one's palm. Such a liquid, showing rapid 
evaporation at room temperature, is said to be volatile and has a 
low boiling-point, e.g., ether, 36°C at standard pressure. 

If a quantity of liquid is contained in an otherwise evacuated, 
sealed glass tube at a certain temperature, molecules will evaporate 
from the liquid surface into the enclosed space. Since these molecules 
are in rapid motion and cannot escape, some of them will collide 
with the liquid surface and re-enter it. It is found that, as a result of 
these opposite processes of evaporation and condensation, a state of 
equilibrium is finally reached, provided some liquid is left. This 
equilibrium is marked by the attainment of a vapour pressure which 
is constant (for the given temperature) and is called the saturated 

vapour pressure of the liquid at that temperature. This saturated 
vapour pressure is independent of the amount of liquid present (pro- 
vided that it is not zero). For example, at 20°C, the saturated vapour 
pressure of water is 17.5 mm. 

Boiling. When a liquid is heated, its molecules acquire increased 
kinetic energy; therefore, the proportion of fast molecules increases, 
evaporation is more rapid and the value of the saturated vapour 
pressure of the liquid rises. For example, for water at 40°C, it is 
55.3 mm and, at 95 C C, 634 mm. When a temperature is reached at 
which the vapour pressure of the liquid equals the prevailing atmos- 
pheric pressure, bubbles of vapour can form freely in the liquid and 
rise to the surface. The liquid is said to boil. That is, the boiling-point 
of a given liquid at a pressure of P mm is the temperature at which 
the vapour pressure of the liquid is P mm. For pure water at 760 mm, 
the boiling-point is 100°C. At a higher pressure, water boils at a 
higher temperature because a higher vapour pressure must be reached 
to achieve free formation of vapour bubbles; conversely, at pressures 
below 760 mm, water has a boiling-point lower than 100°C, e.g., 
110°C at 1075 mm, 90°C at 526 mm. An apparatus for determining 
boiling-points is given on p. 191. 

Other liquids have different boiling-points at standard pressure. 
For example, the binding forces between molecules of cthanol (alco- 
hol) are such that the vapour pressure of this liquid reaches 760 mm 
at 78°C, so this is its boiling-point at standard pressure. The boiling- 
point is higher than 78°C at pressures above 760 mm. 

If a liquid contains non-volatile impurity in solution, this impurity 
will occupy some of the surface of the liquid, hindering vaporisation. 
Consequently, the vapour pressure is lowered at any given tempera- 
ture and the boiling-point of the solution is higher than the boiling- 
point of the pure liquid at the same pressure; correspondingly, the 
freezing-point of the liquid is depressed by non-volatile impurity. For 
example, 3 g of sodium chloride raise the boiling-point of 100 g of 
water by about 0.5°C and depress the freezing-point by about 2°C at 
standard pressure. This is why sea-water is much less readily frozen 
than fresh water in cold weather and why addition of common salt 
causes ice or snow to melt at temperatures not too far below 0°C. An 
experiment illustrating the effect of sodium chloride on the boiling- 
point of water will be found on p. 216. 


Neglecting hydrogen, which has a uniquely simple atom (one pro- 
ton and one electron), the lightest atoms have, in order of atomic 
number from left to right, the following electronic structures. 



He Li Be B C NO F 

Period 2 2 2,1 2,2 2,3 2,4 2,5 2,6 2,7 

Ne Na Mg Al Si P S CI 

Period 3 2,8 2,8,1 2,8,2 2,8,3 2,8,4 2,8,5 2,8,6 2,8,7 

Ar K Ca 

Period 4 2,8,8 2,8,8,12,8,8,2 

It will be seen that, arranged in this way, elements in the same vertical 
columns have the same number of valency electrons in the outermost 
shell of their atoms. Because of this, the elements in each column tend 
to resemble each other closely in chemical behaviour. For example, 
the rare gases. He, Ne and Ar, show a chemical inertness which is 
determined by the stable outer electron octet or duplet. These gases 
form no compounds with other elements. 

Sodium and potassium, each with one electron in the outer shell, 
resemble one another very closely. Both are univalent, ionising by 
the loss of one electron per atom. Both are powerful reducing agents 
and good conductors of electricity. 

Na -> Na + + e~ (K similar) 
Both are strongly electropositive, attacking cold water with liberation 
of hydrogen. 

Na + H t O -»■ Na+OH- + iH, (K similar) 

Both form strongly basic oxides, Na 2 and K 2 0, and soluble hydrox- 
ides, NaOH and KOH, which are strong alkalis. The carbonates, 
Na 2 CO a and K s CO a , are both soluble in water, giving mildly alkaline 
solutions, and are unaffected by ordinary heating. The nitrates of 
both metals liberate oxygen when heated and leave a nitrite. 
2NO a --»-2N<V + 0, 

Lithium resembles sodium and potassium in electropositive char- 
acter and univalency, giving the ion, Li f . It also gives a solid, electro- 
valent hydride, Li + .H - and attacks water liberating hydrogen, but 
only slowly. In this and in its almost insoluble carbonate and fluoride, 
it differs considerably from sodium and potassium and has some 
resemblance to the less electropositive calcium of Group II. 

Similarly, nitrogen and phosphorus show marked chemical similarity 
as non-metals. Both exercise a maximum valency of 5 (i.e., the number 
of outer valency electrons) and also a lower valency of 3 (i.e., the 
octet — 5). These valencies appear in N 2 O s and P 4 O 10 , and in NH 3 
and PH 3 . Both the oxides are strongly acidic and combine with 
water to form acids, HN0 3 and HPO a . 

NjO B + H»0 —*■ 2HN0 3 (P 4 O 10 similar) 



The chlorides, NCl a and PCl a , are both covalent liquids, non- 
electrolytcs and rapidly hydrolysed by water. Both the hydrides, 
NH S and PH a , form salts with hydrogen chloride, NH 4 C1 and 
PH 4 C1, though the latter is much less stable, decomposing at about 

The occurrence of successive groups of elements showing strong 
chemical similarity in this way because of their similar outer electron 
shells is called periodicity. The Periodic Law has been developed fully 
to include all known elements and is expressed in a general Periodic 
Table, of which the above arrangement is the earliest part. Period 1 
(not shown above) is allotted to hydrogen only and Periods 2 and 3 
follow as given on p. 84. The Groups of the table (written vertically) 
arc numbered by Roman numerals, O-VII. See the complete Periodic 
Table printed in the front endpaper. 

The more important of the Groups have already been considered 
briefly for the elements of Periods 2 and 3. Group VII is the halogen 
group; of its members, fluorine and chlorine have been shown in 
Periods 2 and 3 and bromine and iodine follow in later periods. This 
Group is important enough to be given rather special consideration. 
Its members show marked general similarity of properties and also, 
in many instances, a consistent gradation of behaviour. 

Group VII (Halogens) 
These elements and the electronic arrangements in their atoms are: 

F2.7 CI 2,8,7 Br 2,8,18,7 12,8,18,18,7 
Atomic number 9 17 35 53 

The marked similarity of their chemical properties depends es- 
sentially on their possession of the same number (7) of electrons in 
the valency (highest energy) shell, but a pronounced gradation is 
imposed on the general similarity as the atoms increase in com- 
plexity from fluorine to iodine. 

All the elements exhibit a valency of one in covalent combination 
with hydrogen and in electrovalent combination with metals. By 
formation of a shared electron pair with hydrogen in the first case, 
and capture of an electron from a metallic atom in the second, all 
these elements complete the external electron octet. Similarly, by 
formation of one shared electron pair, all produce diatomic mole- 

H, + X 1 ->2HX; X 2 + 2e--^2X-; 2X-+X. 
where X is F, CI, Br or I. 

Being electron acceptors, all the halogens are oxidising agents. 
Examples are the following. 



2Fe a+ + Cl s -*- 2Fe» + + 2C1~ (iron(II) to iron(III)); Br, similar 

S»- + Br a —*■ 2Br- + S (H a S to sulphur); Cl 2 and I 2 similar 

As the number of electron shells increases, the atoms increase in size 
from the smallest, F, to the largest, I. Since each electron layer 
screens the outermost electrons from the attractive power of the 
nucleus, we should expect electrons to be most firmly held (or at- 
tracted) by fluorine atoms and least firmly held (or attracted) by 
iodine atoms. That is, fluorine should be the strongest, and iodine the 
weakest, in oxidising behaviour. This is actually so. For example, in 
the order, F — *■ CI — > Br — >• I, each halogen can oxidise the ions of 
those which follow it and liberate the free halogen. Examples are the 

Chlorine liberates bromine from potassium bromide solution: 
CU + 2Br--*-2Cl- + Br, 

Bromine liberates iodine from potassium iodide solution: 

Br a + 2I--^2Br- + I !1 

There is a similar gradation from F to I in vigour of oxidation of 
hydrogen at room temperature. 

Fluorine combines explosively even in the dark; 
Chlorine combines slowly in daylight; 
Bromine combines slowly in sunlight; 
Iodine combines only when heated and then slowly and 

The same gradation is found in the stability of the hydrides towards 
heat and in the heat of formation of the sodium salts. 

% dissociation 
at 720°C 








Another illustration of the F— >I gradation is given by the 
boiling-points at standard pressure of the four halogen elements, 
which are: 

F 2 




184 (all °C) 

It must be allowed, however, that, in certain respects, halogen pro- 
perties do show some variation. For example, while calcium fluoride 

periodicity 87 

is insoluble in water, the calcium salts of the other three halogens 
(CaCl 8 , CaBr», Cal 2 ) are very soluble and deliquescent. Also, silver 
fluoride is quite soluble in water while the other three silver salts 
(AgCl, AgBr, Agl) have very low solubilities indeed. Also chlorine 
forms an oxide, C1 2 7 , which is the characteristic oxide for a 
Group VII element. None of the other halogens forms such an oxide. 
The highest oxide of iodine is 1 2 6 and of fluorine, F 2 0, while 
bromine forms no stable oxide at all. On strict gradation, bromine 
would be expected to be intermediate between chlorine and iodine in 
solubility in water but, in fact, it is the most soluble of the three (with 
iodine the least soluble). Fluorine attacks water rapidly at room 
temperature, liberating a mixture of ozone, 3 , and oxygen. (For 
some further halogen comparisons relating to Cl 2 , Br 2 and 1 2 only, 
see p. 385.) 

Progression of properties in a Period 

The progression of properties for elements in the same period can 
be illustrated from Period 3. 

Period 3 (Groups in Roman numerals) 


Ne Na Mg Al Si P S CI 

2,8 2,8,1 2,8,2 2,8,3 2,8,4 2,8,5 2,8,6 2,8,7 

The figures represent the electron groupings in the various atoms, 
left to right being outwards from the nucleus. 

The chemical inertness of neon, resulting from its very stable outer 
electron octet, has already been noticed (p. 84). 

Elements of Groups I to III 

Elements. The elements of Groups I, II and III (Na, Mg, Al) all 
show marked metallic character by ionising with electron loss but 
the metallic character weakens in the direction Na — * Mg —*■ Al. 
The valency of each element is equal to the Group number of the 
element and to the number of electrons in the outermost (valency) 

Na-*Na+ + e-; Mg -»- Mg a+ + 2e~; Al -*■ M s+ + 3e~ 

Sodium liberates hydrogen from cold water, showing its exception- 
ally electropositive nature; the other two metals liberate hydrogen 
from dilute acid. In all these cases, hydroxonium ion is reduced by 
electron gain. 

2H.O+ + 2e~ -+• 2H,0 + H, 



Chlorides. The chlorides of sodium and magnesium are electro- 
lytes, Na + Cl~ and Mg 2+ (C1~) 2 . Both are soluble in water; sodium 
chloride is chemically unaffected by water but magnesium chloride 
is hydrolysed slightly, showing the somewhat weaker electropositive 
(metallic) character of magnesium. Aluminium is still less character- 
istically metallic in its chloride. When anhydrous, the chloride is 
covalent (as A1 2 C1 6 ) and is much hydrolysed by water, i.e., resembles 
the chlorides of non-metals, but it yields ions in aqueous solution, 
Al 3+ (hydrated) and CI - as metallic chlorides do. 

Oxides. The oxides of sodium and magnesium are electrovalent 
compounds, (Na + ) 2 2- and Mg 2+ 3- . Both are soluble in water and 
yield alkaline solutions by forming hydroxide ions. The more elec- 
tropositive sodium gives by far the stronger alkali. 

O a " + H 2 -»- 20H~ 

Aluminium oxide has basic properties, e.g., in the reaction 

A1 2 8 + 6HC1 -»- 2A1C1 S + 3H 2 
but it forms no alkali with water and so shows itself less electro- 
positive than sodium or magnesium. Further, aluminium oxide 
shows slight acidic tendency (and is, therefore, amphoteric) by form- 
ing an aluminate with caustic alkali. This resembles the behaviour 
of a non-metal. 

A1 S 8 + 2NaOH + 3H 2 -> 2NaAl(OH) 4 

Hydride. Sodium (heated in dry hydrogen) forms a solid, electro- 
valent hydride. This occurs because the metal is so strongly electro- 
positive as to reduce hydrogen by donating electrons to it. 

2Na + H 2 ->-2(Na + H-) 
This hydride yields hydrogen with cold water and is an electrolyte 
when molten, giving hydrogen at the anode. (Compare hydrides in 
Groups IV to VII later.) 

H~ + H a O -> OH- + H 2 ; 2H~ ->- H 2 + 2er (to anode) 

Elements of Groups IV to VII (Si, P, S, CI) 

In Period 3, all the elements of Groups IV to VII exercise covalency 
numerically equal to their group number by using the valency 
electrons to form covalent pairs, e.g., in oxides SiO s , P 4 O 10 , SO s , 
C1 2 7 . Oxides with lower oxygen content also occur, e.g., P 4 6 , SO z , 
C1 2 0. In addition, each element exercises a covalency of (8 — Group 
number), e.g., in hydrides SiH 4 , PH 3 , H 2 S and HO. Electro-valency is 
shown by elements of Groups VI and VTI only. This is exercised by 
acceptance of electrons to complete the outer octet and so form a 
negative ion. This electrovalency is numerically equal to (8 — Group 


number). Acceptance of more than two electrons per atom never 


S-|-2e--»-S«-; CI + «"-»- Cl~ 

This valency pattern is typical of non-metals. 

Chlorides. The elements of Groups IV and V each produce a 
chloride by utilising covalency numerically equal to the Group num- 
ber, Le., SiCl 4 , PCI B . Both are rapidly hydrolysed by water with 
liberation of hydrogen chloride and, when pure, both are non- 

SiCl 4 + 4H 2 -► Si(OH) 4 + 4HC3; 
PC1 6 + 4^0 -► H 3 P0 4 + 5HC1 

The same elements also produce chlorides by exercising covalency 
of (8 — Group number), Le., SiCI 4 (as before) and PC1 3 . This 
trichloride is a covalent liquid, non-electrolyte and rapidly hydro- 
lysed by water. 

PC1 3 + 3H a O -*■ H 8 PO s + 3HC1 

Contrast the metallic type of chloride produced at the other end of 
the period in Groups I and II, where the elements have 1 or 2 valency 
electrons per atom and produce electrovalent chlorides, solids, 
electrolytes if molten or in aqueous solution, and either unaffected, 
or only slightly hydrolysed, by water. 

Oxides. Elements of Groups IV to VII all form an oxide by exercis- 
ing covalency equal to the Group number, i.e., SiO g , P 4 O 10 , S0 3 
and C1 2 7 . When combined with water, all these oxides produce acids. 
This is characteristic of non-metals. 

Si0 9 + H g O ^ H 2 Si0 3 ; P 4 10 + 2H 2 -> 4HPO„ 
SO a + H 8 -> H 2 S0 4 ; C1 2 7 + H g O -► 2HC10 4 
In addition, phosphorus and sulphur also form well-known oxides 
by exercising a lower covalency, i.e., P 4 0„ and S0 2 . These are also 
acidic oxides. 

P 4 O g + 6H 2 -*■ 4H 8 P0 8 ; SO a + H 2 «* H a S0 8 
Contrast the oxides with basic properties formed by the metals at the 
other end of the Period, Na 2 and MgO. 

Hydrides. Elements of Groups IV to VII all form hydrides by 
utilising covalency equal to (8— Group number), U., SiH 4 , PH 3 , 
H 2 S, HC1. AH these are typical of simple covalent compounds, i.e., 
gaseous at ordinary temperature and pressure and non-electrolytes 
when water-free, having no chemical reaction with water (except 
ionisation by HjS and HCI). Contrast the metallic hydride formed by 
sodium at the other end of the Period -electrovalent (Na + H-), an 
electrolyte when molten (yielding hydrogen at the anode), attacked 
by water to yield hydrogen. 





In physical properties, the Group I-III elements of Period 3 
(Na, Mg, AT) are all metallic, e.g., good conductors of heat and 
electricity. The Group V-VTI elements of Period 3 are non-metallic 
in type, e.g., very poor conductors of heat and electricity. 

Notice the following patterns in Period 3 which are repeated in 
other periods. 


i n in iv v vi vn o 

Valency towards 

ClorH 123 432 1- 

Valency towards 

O (maximum) 12 3 4 5 6 7 — 

Electropositive (metallic) 


character increasing 

Electronegative (non-metallic) 


character increasing 

It is important to know accurately the chemical formula: of the 
chief compounds of the common metals. They are listed below (by 
groups of the Periodic Table) for study and reference. 

Group of P.T. 




Li, Na, K 

Mg, Ca,Ba 








Ca ! + 



(Na+) 2 O a - 

Ca ! +O a - 

(Al 3 +) a (0»-) 3 



Ca'+(OH-) a 

Al 8 +(OH-) 3 



Ca J +(CI-) s 

Al a a 8 (covalent) 


(Na+) 8 S(V- 


(Al s +) a (S(V-) s 



Ca»+(N0 3 -) 3 

Al a +(N0 3 -) 3 


(Na+) 2 CO s "- 










Ca 2+ S 2 - 

(A1»+) 2 (S«-) S 

(decomposed by 


The ionic formula: are shown in most cases though the molecular 
forms are often written, e.g., NaOH, Na 2 S0 4 , Ca(N0 3 ) 2 . 

Ammonium salts have formula; like those of sodium salts, with 
ion, NH 4 + , instead of Na + . There is no oxide. 

Zinc compounds, iron(II) or ferrous compounds, copper(II) or 
cupric compounds and lead(II) or plumbous compounds have ions 
and formulae like those of calcium. 

Iron(III) ox ferric compounds have ions and formula: like those of 

Recent Modifications of Dalton's Atomic Theory 

1 . Elements are made up of small, indivisible particles called atoms 
The atomic nature of elements is not disputed. Atoms can, how- 
ever, no longer be regarded as indivisible in the full meaning of the 
term. Radioactive elements are spontaneously dividing in the sense 
that the atomic nucleus is giving out particles and so producing two 
less complex atoms, e.g., radium disintegrates to produce two rare 
gases, helium and radium emanation (radon). 

2. Atoms of a given element are all exactly alike 

This statement can no longer be accepted. The phenomenon of 
isotopy (p. 77) contradicts it. Thus, potassium has isotopes K = 39 
and K. = 41. Both have 19 nuclear protons and 19 electrons, arranged 

2. 8, 8, 1, so have the same atomic number and properties. But the 
41-isotope has two extra neutrons on the nucleus and so has the 
heavier atom. Most elements exhibit isotopy in this way, (See also 
p. 78.) 

3. Atoms cannot be created or destroyed 

This statement is still acceptable when applied to chemical re- 
actions, in which, apart from electronic changes, atoms react as whole 
units. The changes associated with atomic fission, however, certainly 
destroy atoms of the element involved, in the sense that the nuclei are 
broken into smaller units which correspond to simpler atoms. For ex- 
ample, the nucleus of the uranium isotope, U = 235, can absorb 
a neutron and then break up into two roughly equal parts, which are 
nuclei of atoms in the neighbourhood of barium (Ba = 137). 

4. Atoms combine in small whole numbers 

This statement is still acceptable for most elements. Carbon, 
however, forms the very complex compounds of 'organic' chemistry 
(p.319) and the element, silicon, occurs in some very complex silicates. 


1. What are the three particles which make up ordinary matter? 
Tabulate their relative charges and masses. State briefly how they are 





arranged in a typical form of elementary matter and illustrate your state- 
ment by a diagrammatic representation of one atom of it. 

2. The compounds named below are all covalent compounds. With the 
help of the table of electron-structures on p. 84, give a diagram for a mole- 
cule of each of these compounds, showing the outermost electron shells 
only: (i) carbon tetrachloride, CC1,; (ii) phosphorus trichloride, PC1 3 ; 
(iii) silicomelhane, SiH«; (iv) chloroform, CHC1,; (v) phosphine, PH,; 
(vi) methylene chloride, CH,Cl t . 

3. The following compounds are electrovalent. With the help of the 
table of electron-structures on p. 84, state what electronic changes take 
place when they are formed from their elements: (i) lithium oxide; (ii) 
potassium chloride; (iii) magnesium oxide; (iv) sodium sulphide. State 
briefly what kind of properties you would expect all these compounds to 
show by virtue of their electrovalent character. 

4. Explain the term isoiopy and illustrate by reference to one actual 
example. Explain how this phenomenon contradicts a certain part of 
Dalton's Atomic Theory of 1808. Why, in spite of this contradiction, was 
the Atomic Theory apparently in accord with experimental experience? 

5. Explain briefly the difference between electrovalency and covalency. 
Of which of these two types of valency is ammonia an example? By a 
simple electronic diagram show the essential structure of the ammonia 
molecule. (Give a key to your diagram to explain any symbols used.) 

6. Complete the following table. 
















7. Suppose that elements with chemical symbols. A, B, D, appear in the 
same early Period of the Periodic Table and in Group I, III and V respect- 

(1) Write empirical formula for (a) the basic oxide of A, (b) the sulphate 
of B, (c) the characteristic oxide of D. 

(2) Two of these elements form ions. Write symbols for these ions, 
showing electrical charges. 

(3) Which of these elements is likely to produce an amphoteric oxide? 
Define this term. 

(4) Which of these elements is mostly likely to (a) liberate hydrogen from 
water, (b) form a solid hydride, (c) form a gaseous hydride? Write a 
formula (ionic where appropriate) for each of these hydrides. 

8. (i) In the reaction: Zn + Cu* + — > Zn ,+ + Cu, which reagent is 
oxidised and why? (ii) In the reaction: Cl 2 + 2Br" — >■ 2CI~ + Br„ which 
material is the reducing agent and why? Answer in electronic terms in 
both cases. 

9. Explain the meaning of the term: metallic bond. Explain why a 
typical metal is (a) a good conductor of electricity, (b) a good conductor of 
heat, (c) in some conditions at least, malleable and ductile. 

10. By reference to the properties of (a) the elements, (b) their oxides, 
(c) their chlorides, justify the inclusion of (i) sodium and potassium, (ii) 
nitrogen and phosphorus in the same group of the Periodic Table. 

1 1 . Taking the symbol, l 8 6 X, to represent an atom of the element X, 
state (a) the atomic number of X, (6) the number of neutrons in an atom 
of X, (c) the number of electrons in an atom of X, (d) the mass number of 
X. If another atom is represented as > g 8 X, what term would be used to state 
its relation to ', 6 X and what is the difference between them in terms of 
the number and situation of particles present ? If a sample of X contained 
90% of 'g'X and 10% of ' 8 8 X, show that the chemical atomic weight of X 
would be 16.2. 

12. What feature of the atomic structure of the four halogen elements, 
fluorine, chlorine, bromine, iodine, justifies their common inclusion in 
Group VII of the Periodic Table ? State and explain in electronic terms what 
valency behaviour they all share. Bromine is said to be intermediate in 
behaviour between chlorine and iodine. Justify this statement by reference 
to two examples of the chemical behaviour and one example of the physical 
behaviour of these three elements. Mention one respect in which bromine 
is not the intermediate element of the three and two features of the chemistry 
of fluorine in which it differs from the other three halogen elements. If 
astatine is another element of this Group (following iodine), give one 
chemical property you would expect it to possess and briefly justify your 

13. Explain briefly in electronic terms why (1) the ammonia molecule, 
(2) the oxygen atom can participate readily in co-ordinate bonding. Give 
an example of the formation of an ion from ammonia by this means. Show, 
by electronic diagram, the formation of a co-ordinate linkage between 
phosphorus trichloride and oxygen. State one difference of chemical or 
physical behaviour you would expect between (i) phosphorus trichloride 
and its co-ordinate compound with oxygen, (ii) phosphorus trichloride and 
potassium chloride. Briefly explain your choice. 

14. Suppose that elements W, X, Y, Z are all in the same (early) period 
of the Periodic Table. Allot them to their correct Groups on the following 
evidence: an oxide, W,0 exists and is strongly basic; X forms a liquid, 
covalent chloride, XCIj; the oxide of Y is Y,O a ; Z produces an ion, Z - . 
Answer the following questions: 

(1) What would be the effect of adding W,0 to water and then spotting 
the liquid on to universal indicator paper? Explain. Write an ionic formula 
for the sulphate of W. Briefly characterise the hydride, WH. 

(2) Write the formula of the highest oxide of X and an equation for its 
likely behaviour with water. If X forms a lower oxide, what is its likely 
formula? What is the formula of the simplest hydride of X and its likely 
physical state at s.t.p. ? Relate this to its valency type. 

(3) Write the empirical formula of the chloride of Y. Discuss (briefly) 
likely behaviour of this chloride with water. The oxide, Y,0 3 , is amphoteric. 
Explain this term. 

(4) What is the likely behaviour of the element Z in the field of oxidation- 
reduction? Explain in electronic terms. Write an ionic formula for the 
compound formed between W and Z. Estimate qualitatively its melting- 
point and describe the effect (if any) of an attempt to pass electric current 
through this compound (i) at room temperature, (ii) above its melting- 



15. State, without diagrams, what numbers of electrons occur in the 
successive electron layers of the atoms of (a) carbon, (b) oxygen, (c) nitro- 
gen, (d) chlorine. By diagrams, show the arrangement of valency electrons 
in the molecules of (i) H,0, (ii) ammonia, (iii) carbon dioxide, (iv) 
ethylene, C,H 4 . Sketch the shape of each molecule. Why (in electronic 
terms) is chlorine a posverful oxidising agent? 

16. Show by simple diagrams the arrangement of valency electrons in 
the molecule of (a) methane, (6) acetylene, C»H,. Sketch the shape of each 
of the molecules. Comment on the following facts: (i) the molecules of 
methane and carbon tetrachloride resemble each other very closely in 
shape, (ii) the boiling-points (at standard pressure) of these compounds 
are: methane, 1 12 K, carbon tetrachloride, 350 K. 

Chapter 9 

Molecular Theory 

The Gas Laws 

WE have already seen (Chapter 6) that the behaviour of gases 
when subject to temperature and pressure change, can be ex- 
pressed by two simple laws, those of Boyle and Charles. 
Boyle's Law. The volume of a given mass of gas is inversely pro- 
portional to its pressure, temperature remaining constant. 
With the usual symbols, this is expressed mathematically as: 

pv = a constant (T constant) 
Charles' Law. The volume of a given mass of gas is directly pro- 
portional to its absolute temperature, pressure remaining constant. 

With the usual symbols, this is expressed mathematically as: 

— = a constant (p constant) 

Gay-Lussac's Law of gaseous volumes 

A third law, describing the behaviour of gases, when involved in 
chemical reactions, was stated by Gay-Lussac. 

We can illustrate the Law of Gay-Lussac by quoting first some of 
the experimentally observed results of chemical reaction between 
gases, upon which the law is based. Temperature and pressure are 
to be considered constant throughout each statement. 

1. Ammonia. 

2 volumes of ammonia decompose to give 1 volume of nitrogen 
and 3 volumes of hydrogen. 

2. Steam. 

2 volumes of hydrogen combine with 1 volume of oxygen, giving 
2 volumes of steam. 

3. Hydrogen chloride. 

1 volume of hydrogen combines with 1 volume of chlorine to 
give 2 volumes of hydrogen chloride. 




4. Nitric oxide {nitrogen monoxide) 

2 volumes of nitric oxide decompose to give 1 volume of nitro- 
gen and 1 volume of oxygen. 

Examining these experimental results (the methods by which they 
have been obtained are given in Chapter 1 1), we notice at once that 
all the volumes of the gases concerned are related to each other by 
simple whole number ratios. 

Whenever gases are concerned in chemical action, simple whole- 
number relations between their volumes are always found. This is the 
fact which was first noted by Gay-Lussac and expressed in his Law of 
Gaseous Volumes, which is now stated. 

Gay-Lussac's Law of gaseous volumes. When gases react they do so in 
volumes which bear a simple ratio to one another, and to the volume of 
the product if gaseous, temperature and pressure remaining constant. 

Simple behaviour of gases: an explanation required 

These three Laws of Boyle, Charles and Gay-Lussac express among 
them a highly interesting fact about gases — a curious similarity of 
behaviour. In chemical properties, and such physical properties as 
density and solubility in water, gases show marked variations. There 
are neutral gases such as nitrogen, oxygen and hydrogen, acid- 
producing gases such as sulphur dioxide, hydrogen chloride and 
nitrogen dioxide, alkali-producing gases such as ammonia; gases of 
very high solubility in water, such as hydrogen chloride (500 volumes 
of gas dissolve in 1 volume of water), gases of moderate solubility, 
such as hydrogen sulphide (3 volumes of gas dissolve in 1 volume of 
water) and gases of very low solubility, such as nitrogen (0.02 volumes 
dissolve in 1 volume of water); some gases are chemically very 
reactive, e.g., chlorine, and some are entirely inert, e.g., argon. But, 
however great the variations in these properties may be, all the gases 
obey the Laws of Boyle, Charles, Gay-Lussac. There must be some 
explanation of this similarity. Note that it does not matter whether 
the gas is an element, e.g., hydrogen, or a compound, e.g., hydrogen 
chloride; each obeys the laws equally well. 

Avogadro's Hypothesis 

The explanation was put forward in 181 1 by Avogadro, an Italian 
scientist, in the form known as Avogadro's Hypothesis. We have 
seen in Chapter 2 that the smallest particle of an element or com- 
pound which can exist separately is called a molecule of it. Avogadro's 
explanation of the simple behaviour of gases, especially as expressed 
in Gay-Lussac's Law, was that equal volumes of all gases, under the 
same temperature and pressure conditions, contain the same number 
of molecules. When this suggestion was put forward it was purely a 



hypothesis, that is, an idea which had occurred to Avogadro, which 
appeared to him sensible, but which still required to be tested further 
before it could be fully accepted. The truth of it has since become 
experimentally demonstrable, and it is frequently known, on that 
account, as Avogadro's Law. (See Avogadro Constant, p. 107.) 
Avogadro's Law. Equal volumes of all gases at the same temperature 
and pressure contain the same number of molecules. 

This law has been of the greatest value in the development of 
chemistry since about 1860. It is a rather curious fact that its im- 
portance was at first unnoticed, and the full recognition of its im- 
plications, a few of which we shall now examine, is due to the work, 
not of Avogadro himself, but of another Italian, Cannizzaro, some 
47 years after the hypothesis had been first put forward, and after 
Avogadro himself was dead. 

Why Avogadro's Law is important 

The importance of the law lies in this fairly simple fact, 
that, since it asserts that equal volumes of gases contain 
equal numbers of molecules, it enables us to change over 
directly from a statement about volumes of gases to the same 
statement about molecules of gases. Every time we make a 
statement about one volume of any gas, we are also making 
a statement about a certain number of molecules of it, and 
that number, by Avogadro's law, is always the same, no 
matter what the gas may be. Consequently, we can change 
over at will, in any statement about gases, from volumes to 
molecules and vice versa. 
This means that by applying the law to volume measurements of 
gases, we can probe right to the heart of a chemical reaction, to the 
actual molecules themselves. It is an enormous step to change 
directly from an experimental statement like: 
2 volumes of hydrogen combine with 1 volume of oxygen giving 
2 volumes of steam (temperature and pressure constant) 

2 molecules of hydrogen combine with 1 molecule of oxygen giving 

2 molecules of steam. 
The second of these two statements goes right to the essentials of 
the reaction, to the very molecules themselves. The law is important 
because it gives us this power to reveal the molecules themselves at 
work in chemical reactions. Note, however, that it applies only to 

Other examples illustrating this important change from volume 
measurements to statements about molecules follow overleaf. 





Ammonia. By experiment, 2 volumes of ammonia decompose to 
give 1 volume of nitrogen and 3 volumes of hydrogen. 

Applying the Law, 

2 molecules of ammonia contain 1 molecule of nitrogen and 

3 molecules of hydrogen. 

Hydrogen Chloride. By experiment, 1 volume of hydrogen com- 
bines with 1 volume of chlorine to give 2 volumes of hydrogen 

Applying the Law, 

1 molecule of hydrogen combines with 1 molecule of chlorine 
to give 2 molecules of hydrogen chloride. 

Nitric Oxide. By experiment, 2 volumes of nitric oxide decompose 
to give 1 volume of nitrogen and 1 volume of oxygen. 

Applying the Law, 

2 molecules of nitric oxide contain 1 molecule of nitrogen and 
1 molecule of oxygen. 

(Temperature and pressure assumed constant throughout) 

From these statements, it is only a further step to the deduction of 
the formula of the gaseous compounds concerned. This step is given 
in the later part of this chapter, and the formulas of the common 
gases are considered in the next. 

How Avogadro's Law explains Gay-Lussac's Law 

Assume throughout the following paragraph that temperature and 
pressure are constant. 

When gases react chemically, the reaction must take place between 
individual molecules of the gases. As Dalton suggested in the similar 
case of combination between atoms, the reactions will take place 
between small whole numbers of molecules of the reactants to pro- 
duce small whole numbers of molecules of the products. 

We have seen in the last section that, employing Avogadro's Law, 
we can change over directly from statements about molecules to 
statements about volumes, provided that gases only are concerned. 
Making this change, the last sentence of the last paragraph becomes: 
the reactions will take place between small whole numbers of volumes 
of the reactants to produce small whole numbers of volumes of the 
products (all being gases). This is what Gay-Lussac's Law states. 
Hence, Avogadro's Law has enabled us to deduce the experimentally 
observed Law of Gay-Lussac. 

Avogadro's Law and the molecular weights of gases 

We have seen, in Chapter 4, that molecular weights are expressed 
as the number of times one molecule of the substance is as heavy as 
one atom of hydrogen. We have also seen, in this chapter, the very 
important relation which exists between the number of volumes of 
gases and the number of molecules of gases involved in chemical re- 
action. It is now necessary to find how the atom and the molecule of 
hydrogen are related to one another. This will lead us to a method 
of determining molecular weights. 

The nature of the hydrogen molecule 

By experimental work which is fully described later (p. 1 12), it has 
been found that (at constant temperature and pressure): 

1 volume of hydrogen combines with 1 volume of chlorine to give 
2 volumes of hydrogen chloride. 

Applying Avogadro's Law we can say at once: 

1 molecule* of hydrogen combines with 1 molecule of chlorine to 
give 2 molecules of hydrogen chloride. 

Now each of the two molecules of hydrogen chloride must 
contain some hydrogen. The least amount of hydrogen 
which can be contained in one molecule of hydrogen 
chloride is one atom, because the atom of hydrogen is in- 
divisible. Consequently, the least amount of hydrogen there 
can be in two molecules of hydrogen chloride is 2 atoms of 
hydrogen. But these 2 atoms of hydrogen must have come 
from the 1 molecule of hydrogen marked*; therefore, a 
molecule of hydrogen must contain at least two atoms of 
But hydrochloric acid forms with sodium hydroxide one salt only, 
sodium chloride. Two series of salts have never been obtained from 
hydrochloric acid as they have from, for example, sulphuric acid. 
Thus with sodium hydroxide, sulphuric acid can be made to form 
both normal sodium sulphate and acid sodium sulphate. Since the 
hydrogen of hydrochloric acid cannot be replaced in two stages, 
there is only one hydrogen atom in the molecule. But the two 
hydrogen atoms necessary for two molecules of hydrogen chloride 
have been obtained from one molecule of hydrogen. Hence the 
molecule of hydrogen contains two atoms. 

Try to visualise what this means. It means that, in ordinary gaseous 
hydrogen, no separate atoms of hydrogen exist. All the particles 
consist of two hydrogen atoms, locked in a chemical embrace, and 



moving always as a single unit, the molecule. It is as if the hydrogen 
atoms are paired off to run a perpetual three-legged race. The mole- 
cule, consisting of two hydrogen atoms, never breaks up, except for 
the purpose of engaging in chemical reactions. This fact is expressed 
by writing the hydrogen molecule as H t , which means a single mole- 
cule of hydrogen containing two atoms. (2H would mean two separ- 
ate hydrogen atoms.) 

By a somewhat similar argument, it can be shown that the mole- 
cule of chlorine contains two atoms, and the reaction between 
hydrogen and chlorine may be diagrammatical ly expressed as: 







ft 8 8 ft 

ft ft * *, 

ft <* ft * 

1 volume 

■ 1 volume . 

+ 6C1, "* 

2 volumes 



or in simplest terms, 


+ CI, -► 
Fig. 30. 


Note that equal numbers of molecules of hydrogen, chlorine and 
hydrogen chloride are contained in equal volumes. 

It is known, from similar experimental evidence and argument, that 
nitrogen and oxygen also have two atoms per molecule and their 
molecules are written N s and 2 . This is expressed by saying that 
the molecules of hydrogen, chlorine, nitrogen and oxygen are 
diatomic or that their atomicity is 2. 

Atomicity. The atomicity of an element is the number of atoms con- 
tained in one molecule of the element. 

Relation between vapour density and molecular weight 

The relative densities of solids and liquids are expressed with 
reference to water, but it would be most inconvenient to deal with 
gases in this way because of the great difference between the densities 
of water and gases. 

Definition. The vapour density of a gas or vapour is expressed as the 
number of times a certain volume of the gas or vapour is as heavy as 
the same volume of hydrogen at the same temperature and pressure. 
Expressed in another form, this becomes: 

Vapour density of _ Weight of 1 volume of gas or vapour 
a gas or vapour We i ght of i volume of hydrogen 

Temperature and Pressure constant 



Note that vapour density can be experimentally determined 
because it only involves weighing equal volumes of hydrogen and 
the vapour. 

The molecular weight of a gas or vapour is expressed in the lorm: 

Molecular weight of _ Weight of 1 molecule of the gas or vapour 
a gas or vapour Weight of 1 atom of hydrogen 

We shall now show that there is a simple relation between vapour 
density and molecular weight. 

Vapour density of _ Weight of 1 volume of gas or vapour 
a gas or vapour Weight of 1 volume of hydrogen 

Temperature and Pressure constant 
Applying Avogadro's Law, we can say directly: 
Vapour density of _ Weight of 1 molecule of gas or vapour 
a gas or vapour Weight of 1 molecule of hydrogen 

_ Weight of 1 molecule of gas or vapour 
Weight of 2 atoms of hydrogen 
Multiplying both sides by 2: 
2 X (Vapour Weight of 1 m olecule of gas or vapour 

density of a gas = Weight of j atom of hydrogen 

or vapour) 

= Molecular weight of the gas or vapour 

/.c, the molecular weight of a gas or vapour is twice its vapour 

Molecular weight from vapour density. Regnault's Method 

We have already noted that to find the vapour density of a gas or 
vapour it is only necessary to obtain the weight of a certain volume of 
the gas or vapour and the weight of an equal volume of hydrogen, 
both at the same temperature and pressure. 

Unfortunately, direct weighing of hydrogen and other gases in this 
way is very difficult, partly because the actual weights of convenient 
volumes of the gases are small, and partly because changes of tem- 
perature, pressure and humidity in the atmosphere introduce errors 
during the course of the experiments. 

In principle it is only necessary to evacuate a globe, weigh it and 
fill it with hydrogen and weigh it; then evacuate it again, fill with the 
gas and weigh again, temperature and pressure remaining constant. 
This is known as Regnault's Method. Then, 
Vapour density _ (Weight of globe + gas) — (weight of globe) 
of the gas ~ (Weight of globe + hydrogen) — (weight of globe) 

and Molecular weight of the gas = 2 X vapour density. 



A simpler method of finding molecular weight, which can be 
applied to the particular case of oxygen, is described in the next 

Determination of molecular weight of oxygen 
This method is simplified by: 

1. Weighing the oxygen as a loss in weight of potassium chlorate. 

2. Using the fact that 1 dm 3 of hydrogen at s.t.p. weighs 0.09 g, 
and measuring the volume of oxygen. 

Weigh a hard glass test-tube containing some potassium chlorate. 
Use apparatus of Fig. 11, p. 46. Fill the siphon tube with water by 
blowing, with clip open, into the short aspirator tube. Close the clip. 
Attach the hard glass test-tube to the short aspirator tube and put 
the measuring cylinder into position. Open the clip. (Water will flow 
for a second or two; longer flow indicates leakage which must be 
corrected.) Heat the potassium chlorate. Oxygen evolved will displace 
water into the measuring cylinder. When it is suitably filled, let the 
apparatus cool. Equalise water-levels in measuring cylinder and 
aspirator (so that oxygen is at atmospheric pressure). Close the clip, 
remove the measuring cylinder and read the volume of water (i.e., 
of oxygen). Weigh the hard glass tube and record room temperature 
and pressure. 


Weight of hard glass tube and potassium chlorate 21.32 g 
Weight of hard glass tube after heating 21.03 g 

.". Weight of oxygen = 0.29 g 

Volume of oxygen = 217 cm 3 

Temperature = 15°C 

Pressure (corrected for vapour pressure of water) = 750 mm 

773 7Sfl 

Volume of gas at s.t.p. = 217 x =^ x i£ cm 3 

288 760 

= 203 cm s 
.'. 1000 cm 3 of oxygen weigh 029 * 100 ° g = 1.43 g at s.t.p. 
But 1000 cm 3 of hydrogen at s.t.p. weighs 0.09 g 

/. Vapour density of oxygen = -^— = 15.9 

.\ Molecular weight of oxygen = 15.9 X 2 = 31.8 

Scales of atomic weight 

It was stated in Chapter 4 that, after the mid-nineteenth century, 
the atomic weight scale was based on the reference standard, H = 1. 



On this standard, the atomic weight of oxygen was 15.88, or 
O = 15.88. This atomic weight scale continued in use until the early 
years of the twentieth century when it was found desirable to make a 
slight modification by basing the scale on the standard of O = 16. 
This meant multiplying all existing atomic weights by the fraction 

J^L so that, for example, the atomic weight of hydrogen was raised 


to 1.008. 

The principal reason for making this change was that many 
equivalents of elements (and hence their atomic weights) were 
determined by a direct experimental relation with oxygen. If H = 1, 
hence O = 15.88, was used as standard, any error in the figure 15.88 
was transferred to the equivalents of these many elements. If, how- 
ever, the standard was directly O - 16, hence H = 1.008, any error 
in the 1.008 affected hydrogen only. In conformity with this new 
scale, the definition of equivalent of an element (p. 37) was modified 
to relate to 8 mass units of oxygen exactly and to the corresponding 
figures for the other elements used experimentally. 
Definition. The gram-equivalent weight of an element is the number of 
grams of the element which combine with or displace 8 grams of 
oxygen, 1.008 grams of hydrogen or 35.46 grams of chlorine. 

This scale continued in use until January 1, 1962, when a new scale 
based upon carbon was substituted. 

There were two reasons for this change. The first was that, about 
1930 it was found that ordinary atmospheric oxygen is not a uniform 
material but contains three isotopes, l 8 O (99.76%), ' 8 7 (0.04%) and 
^O (0.02%). If the atomic weight of oxygen-16 is taken as O = 16, 
the average atomic weight of atmospheric oxygen is 16.0044. Further, 
the isolopic content can vary slightly so that chemical oxygen does 
not provide a constant standard. The second reason for the change 
to carbon as reference element was that, by the 1950's, the principal 
means for atomic weight determination had become the mass spectro- 
meter to which carbon is much more suited than oxygen. 

So, after international consultations between chemists and physi- 
cists in 1960-1, it was agreed to adopt an atomic weight scale based 
upon the carbon isotope, « *C, as C - 12. (This isotope is usually 
referred to as carbon-12 or ,2 C.) The numerical effect on atomic 
weights was only slight— a reduction of about 37 parts per million 
(e.g., O - 16.0000, became O - 15.9994). The chemical atomic 
weight of carbon is not exactly 12; the presence of a small proportion 
of the isotope, l 3 C, raises it to 12.01115, with an error of ±0-0°™» 
because of variation in isotopic composition. The reference standard 
for equivalents must now be 3 mass units of carbon (since the valency 
of carbon is 4) but the former reference standards of 8 units of oxygen, 



1.008 units of hydrogen and 35.46 units of chlorine remain numeri- 
cally unaffected at the ordinary standards of experimental accuracy. 

Molar volume (or gram-molecular volume) of hydrogen 

It was shown on p. 99 that the hydrogen molecule is diatomic and 
is written as H 2 . Expressing this on the standard of ia C = 12 and 
H = 1.008, we have: 


2 X 1.008 
or 2.016 

If this is expressed in the scientific weight unit (gram), it becomes 
2.016 g and this is called one mole of molecules of the gas (or one 
gram-molecule of it). It contains the Avogadro Constant (or Number) 
of molecules, 6.02 x 10 23 . 

Definition. The molar weight (or gram-molecular weight, G.M.W.) of 
any gas is its molecular weight expressed in grams. 
We now have : 

Molar weight 2.016 g (or G.M.W.) 

By experiment, it has been found that 1 dm 3 of hydrogen at s.t.p. 

weighs 0.09 g. Therefore, 2.016 g of hydrogen occupy 


dm 8 , or 

22.4 dm 3 , at s.t.p. This volume is called the molar volume (or gram- 
molecular volume) of hydrogen. 

We see, from this, that if we use grams as our weight-unit, the 
formula, H s , may denote either 2.016 g of hydrogen or 22.4 dm 3 of it 
at s.t.p. We have connected the molecular formula of hydrogen with 
a volume, rather a convenient result because hydrogen, a gas, is 
usually measured experimentally as a volume. 

Molar volumes (or gram-molecular volumes) of other gases 

Consider the same volume, 22.4 dm 3 at s.t.p., of some other gas, 
say oxygen. Since we are considering the same volume of both oxygen 
and hydrogen in the same conditions, we know, by Avogadro's Law, 
that we must be considering the same number of molecules of the two 
gases. But we started from one mole of molecules of hydrogen; there- 
fore, the 22.4 dm 3 at s.t.p. must represent the same number of mole- 
cules of oxygen, i.e., one mole of molecules (or one gram-molecule) 
of it. The same argument will apply to any other gas, which gives us 
this very important result: 
One mole of molecules (or one gram-molecule) of any gas occupies 



22.4 dm 3 at s.t.p. This volume is called the molar volume (or gram- 
molecular volume) of the gas. 

Definition. The molar volume (or gram-molecular volume, G.M.V.) of 
any gas is the volume occupied at s.t.p. by one mole of molecules (gram- 
molecule) of the gas and is 22.4 dm 3 . 

This is a most important and useful result because it means that if 
we write the molecular formula of any gas and refer it to grams as 
weight-units, the volume of gas indicated is always 22.4 dm 3 at s.t.p. 


H 2 O, N 2 C0 2 H 2 S SO s CI* 
2.016 32 28 44 34 64 71 grams 
22.4 22.4 22.4 22.4 22.4 22.4 22.4 dm 3 at s.t.p. 

This means that the gases all have differing densities except where, by 
coincidence, their molecular weights are the same, e.g., CO» and 
NA 44. 

Application of molar (gram-molecular) volume of gases to determina- 
tion of molecular weights 

It follows from the above result that, to determine the molecular 
weight of a gas in grams, we have simply to find the weight of the gas 
in grams which occupies 22.4 dm 3 at s.t.p. 

Example. 350 cm 3 of a certain gas were found to weigh 1 g at s.t.p. 

conditions. What is the molecular weight of the gas? 

From the data given, the weight of 22.4 dm 3 (or 22 400 cm 3 ) of the 

I x 22400 
gas at s.t.p. is — grams or 64 grams. 

Therefore, the molecular weight of the gas is 64. This method of 
calculation is alternative to the vapour density method given on 
p. 101. 

Application of molar (gram-molecular) volume to gas calculations 

Example 1. Calculate the volume of oxygen at 12°C and 745 mm 
pressure which could be obtained by heating 5 g of potassium chlorate. 
(K = 39; CI = 35.5; 0=16. Molar, or gram-molecular, volume of 
gases at s.t.p. is 22.4 dm*.) 

Calculation. The first requirement is the equation: 
2KC10 3 -*■ 2K.C1 + 30 8 
Note that, in the question, a weight of potassium chlorate is given and 
a volume of oxygen is wanted. Therefore, in the equation, we insert 
the weights appropriate to potassium chlorate and the volume appro- 
priate to oxygen. It is quite unnecessary to insert a weight of oxygen. 



Using grams as the weight units for potassium chlorate, O, (one 
mole, or gram-molecule, of oxygen) represents 22.4 dm 3 at s.t.p.; 
30 a represents, therefore, 3 x 22.4 dm 3 , or 67.2 dm». 
Inserting weights and volume, we have: 

2KC10 8 -> 2KC1 + 30 8 

2(39 + 35.5 + 48) g 67.2 dm 3 

245 g at s.t.p. 

From the equation: 

245 g of potassium chlorate yield 67.2 dm 3 of oxygen, 

so 5 g of potassium chlorate yield — ' _ * - dm 3 of oxygen, 

= 1.37 dm 3 of oxygen at s.t.p. 

Converting this volume to 12°C and 745 mm pressure as the example 

Pi»i _ Pi"i 

f t r, 

760 X 1.37 745 X v t 


"» = 

760 x 1.37 X 285 

273 x 745 
= 1.46 dm* 

dm 3 

Example 2. Calculate the volume of hydrogen sulphide at 14° C and 
110 mm pressure which will react with 10 g of lead(IT) nitrate. 
(Pb = 207; N = 14; O = 16. Molar, or gram-molecular, volume of 
gases is 22.4 dm 3 at s.t.p.) 

Calculation. We require first the equation; the nitrate is given as a 
weight so we insert weights under its formula, while the hydrogen 
sulphide is required as a volume so the volume appropriate to one 
mole of gas is put under its formula. 

H t S + Pb(NO,), ->-PbS + 2HNO a 

22.4 dm 3 (207 + 28 + 96) g 
at s.t.p. 331 g 

From the equation : 

331 g of the nitrate react with 22.4 dm 3 of hydrogen sulphide, 

22 4 x 10 
so 10 g of the nitrate react with — '-—— dm 3 of hydrogen 

sulphide at s.t.p. = 0.677 dm 3 at s.t.p. 


Converting this volume to 14°C and 770 mm as the example 

PlVl = PjP* 

760 x 0.677 770 X P, 


Of = 

760 X 0.677 X 287 

273 X 770 
= 0.703 dm 3 

dm 3 

Avogadro Constant 

The Avogadro constant has the value, 6.02 X 10* 3 and is defined 
as the number of atoms in exactly 12 g of carbon-12. It is often 
called the Avogadro number. It is also the number of molecules 
contained in one mole of molecules of any gas. It is, for example, 
the number of molecules contained in 36.5 g of hydrogen chloride, 
32 g of oxygen or 64 g of sulphur dioxide, these figures being the 
molecular weights of the gases. The volume of gas which contains 
this number of molecules always occupies 22.4 dm 3 at s.t.p. 

The Avogadro number is also the number of atoms contained in 
a mole of atoms of any element, e.g., in 22.99 g of sodium or 26.98 g 
of aluminium, the figures being the atomic weights of the elements. 
An older name for a mole of atoms was one gram-atom, i.e., the 
atomic weight of the element in grams. The Avogadro number of 
electrons is also called a mole of electrons and can liberate in electro- 
lysis one mole of atoms of any univalent metal. 

An early method for the determination of the Avogadro number 
(1912) came out of the work of an American, Millikan, who deter- 
mined the charge on the electron as 1.59 x 10~ 19 coulomb. Since 
96 500 coulombs of electricity were known to liberate one gram-atom 
of any univalent element in electrolysis (e.g., 1.008 g of hydrogen), 

the number of atoms involved must be - ^_ u or 6.07 X 10 23 . 

As stated above, the accepted value today is 6.02 x 10 18 (accurate 
to 3 sig. fig.). 

Empirical and molecular formula; 

A method was given on p. 56 for the calculation of a formula for a com- 
pound from its percentage composition by weight. It is quite possible for 
two different compounds to have the same percentage composition by 
weight either because the compounds have different arrangements of the 
same atoms inside the molecule, or because the molecular formula of one 



is a multiple of that of the other. Considering the second of these possi- 
bilities, it is clear that acetylene, C a H„ and benzene, C,H 6 , both having 
92.3% of carbon, will both appear to have the same formula, when the 
calculation of p. 56 is applied. Thus: 

Per cent by weight 

Number of atoms is represented by 

Dividing by smallest 



12 '•' 

^ 7 =1 



The formula appears to be CH for both. The reason is that this cal- 
culation always yields the simplest formula which expresses the com- 
position of the substance by weight. Since the ratio of carbon atoms to 
hydrogen atoms is the same in CH, QH, and C,H„ the same composition 
by weight is expressed in all three. This simplest formula which expresses 
the composition of a compound by weight is called its empirical formula. 
Thus, the empirical formula of both benzene and acetylene is CH. 

Clearly we must devise a method of finding the true or molecular 
formula of the compounds. This is simple enough. If the true formula is 
CH, the molecular weight is (12 + l)or 13;ifC,H„26;irc,H„78,andso 
on. Thus, a determination of molecular weight will at once decide the true 
formula, and, in practice, this means determining the vapour density 
(p. 101) of the compound. The vapour density of acetylene is 13, and of 
benzene 39; that is, their molecular weights are 26 and 78 respectively. 
This gives a molecular formula C,H a for acetylene and QH, for benzene. 

Definitions. The empirical formula of a compound is the simplest 
formula which expresses its composition by weight. 

The molecular formula of a compound is one which expresses the 
actual number of each kind of atom present in its molecule. 

Another example will illustrate the point further. 

A gaseous compound of carbon and hydrogen contains 80% carbon by 
weight. One dm' of the compound at s.t.p. weighs \.35g. Find its molecular 
formula. (C = 12; H - 1. G.M.V. of any gas is 22.4 dm* at s.t.p.) 

Per cent by weight 

Number of atoms is represented by 

Divide by smallest 


The empirical formula is CH, 

.-. The molecular formula is CH,,, where n is a whole number, 

.-. the molecular weight is (12/i + 3/i). 

From the problem, 1 dm 3 of the compound at s.t.p. weighs 1.35 g. 

. . 1.35 x 22.4 ,«,„ 
.-. 22.4 dm s of the compound at s.t.p. weigh j — =30.2 g 

12n + 3/» = 30.2 
I5« = 30.2 
n = 2 
.-. Molecular formula is C,H, (ethane) 

Application of Avogadro's Law to determination of atomic weights 

We have seen in Chapter 4 the great difficulty encountered, in the first 
half of the nineteenth century, in the determination of atomic weights and 
how the difficulty was overcome in the case of metals by the use of equiva- 
lent and valency. A development from Avogadro's Law, which we are 
now to consider, supplied a method of determination of the atomic weight 
of non-metals of a certain type. 

The most convenient case for us to consider is that of carbon. Suppose 
we take the symbol X to denote the atomic weight of carbon. In a molecule 
of a carbon compound, there cannot be less than one atom of carbon, 
and there may be two, three, four or any small whole number of carbon 
atoms. This means that, in the molecular weight of a carbon compound, 
there must be X, 2X, 3X or n X units of weight of carbon (n is a small whole 

number). , 

The molecular weight of any carbon compound can be found by deter- 
mining first its vapour density (by the method on p. 101, provided that the 
compound is gaseous). The compound can then be analysed and the 
percentage of weight of carbon in it determined. The weight of carbon in 
the molecular weight is then given by the expression. 

Percentage of c arbon Mol ecular weight of compound 

TOO 1 

If this is applied to several carbon compounds, the results must represent 
the weights of the number of carbon atoms in the molecules. In the table 
on p. 110, the figures are given for several compounds. 

The figures in the last column correspond to the presence or one, two, 
three or more carbon atoms. The lowest weight is 12 and the others are 
multiples of 12. Now it is obvious that, if we have included in our list any 
compound containing only one carbon atom per molecule, that com- 
pound will be the first, methane, or the last, formaldehyde, because in these 
the weight of carbon is the least. If, therefore, the molecules of methane 
and formaldehyde do actually contain only one carbon atom, the atomic 
weight of carbon is 12. This process has been applied to a very large 
number of carbon compounds, and the weight of carbon in the molecular 
weight has always been found to be 12, or a multiple of 12, but never less. 
From this we conclude that the least weight of carbon there can ever be in 
the molecular weight of one of its compounds is 12, that this weight 
corresponds to the presence of one carbon atom and that the atomic weight 
of carbon is 12. 



The method can be applied to determine the atomic weight of any 
clement forming a large number of gaseous or easily volatile compounds. 


density by 



% of carbon 
by (weight by 

Weight of carbon 

in the molecular 











75 x j£ = i2 
80 x-^24 
81.8 x « - 36 
85.7x^ = 24 

40 »*-« 


All necessary atomic weights may be found on p. 532. 

1. What is the weight of 22.4 dm 3 (gram-molecular volume) of the 
following gases at s.t.p. (a) ammonia; (6) hydrogen sulphide; (c) nitro- 
gen; (a 1 ) chlorine; (e) dinitrogen oxide? 

2. Calculate the molecular weights of the following gases from the state- 

(a) 0.8 g of oxygen occupied at s.t.p. a volume of 560 cm*. 

(b) 1400 cm' of sulphur dioxide measured at s.t.p. weighed 4 g. 

(c) 1.12 dm' of nitrogen monoxide measured at s.t.p. weighed 1.5 g. 

3. What volume of carbon dioxide at s.t.p. could be obtained by 
dissolving 1 50 g of pure marble (calcium carbonate) in dilute nitric acid ? 

4. 1 . 1 6 g of magnesium was allowed to react with excess dilute sulphuric 
acid. What volume of hydrogen measured at s.t.p. was liberated? 

5. An evacuated flask weighed 20.70 g. Filled with dry hydrogen, it was 
found to weigh 20.94 g. Filled with dry chlorine at the same temperature 
and pressure as the hydrogen, the flask weighed 29.22 g. Using these data 
alone, find the molecular weight of chlorine. 

6. 1 dm* of ozone measured at 20°C and 750 mm was converted into 
oxygen by heating. If the resulting oxygen was measured at 30°C and 
750 mm, what volume would it occupy? 

7. 100 cm' of hydrogen were sparked with 30 cm' of oxygen, both gases 
at 1 10°C and 760 mm. What is the total volume of gas left after cooling to 
the original temperature and pressure? What percentage of this gas by 
volume is steam? 



8. By heating a certain weight of lead dioxide with excess of concen- 
trated hydrochloric acid, 743 cm 3 of chlorine were obtained, measured at 
27°C and 755 mm. Calculate the weight of lead dioxide used and find 
the weight of lead chloride which was produced during the reaction. 

9. Explain carefully the term 'molecular weight'. Give a short account 
of the experiments and reasoning which lead to the conclusion that the 
molecular weight of steam is 18. (O. and C.) 

10. Define the terms 'vapour density* and 'molecular weight', and 
deduce from your definition the relation that exists between them, stating 
any assumption that you make in the deduction. A metallic chloride has a 
vapour density of 130 and contains 54.6% of chlorine by weight. How 
many atoms of chlorine does its molecule contain? (CI = 35.5.) (C.) 

1 1. What is the relation between molecular weight and vapour density, 
and how do you account for it? 

At atmospheric pressure and 546°C, 50 cm 3 of phosphorus vapour weigh 
0.093 g. What is the molecular weight of phosphorus (1 dm 3 of hydrogen 
at s.t.p. weighs 0.09 g)? (O. and C.) 

12. Explain how Avogadro's Law can be used to establish (a) molecular 
weights; (o) atomic weights. Illustrate your answer by reference to oxygen 
and carbon. (O. and C.) 

13. A flask of about 500 cm 3 capacity, fitted with a rubber stopper, 
weighed 90.512 g when filled with air. Hydrogen chloride from a generator 
was passed in for a few minutes, the stopper was replaced, and the flask 
and its contents now weighed 90.660 g. 

The flask was inverted in a trough of water, the stopper was removed 
under water, and eventually 456 cm 3 of liquid entered when the levels had 
been adjusted. Given that the temperature of the gases was 13 C and that 
the atmospheric pressure was 770 mm of mercury, find: 
(a) The weight of a litre of hydrogen chloride at s.t.p. 
(£) The molecular weight of hydrogen chloride from the given data. 
(One dm 3 of hydrogen at s.t.p. weighs 0.09 g, one dm 3 of air at 
s.t.p. weighs 1.293 g.) (B.) 

14. Explain the use of a chemical equation and the information which 
it conveys. 

Give equations representing three chemical reactions and state the exact 
meaning of each equation. (O. and C.) 

15. Describe in detail the qualitative and quantitative experiments you 
would carry out to show that the equation CuO + H, = H a O + Cu re- 
presents the action of hydrogen on red-hot copper oxide. (O.) 

Chapter 10 

Formula of Gases 

It is important to mention the conditions of temperature 
and pressure whenever gas volumes are measured 

TN this chapter, we shall give the experimental evidence, and the 
Areasonmg from it, by which the formula? of a number of common 
gases have been established. It is most important that you should 
note carefully the way in which Avogadro's Law is continually 
employed in establishing these formula;, both directly in converting 
volume measurements into evidence of the numbers of molecules 
involved and, indirectly, when vapour density measurements are 
employed. The importance of the Law in the following work cannot 
be over-estimated. 

Hydrogen chloride 

All volume measurements made during the following experiment 
are under the same conditions of temperature and pressure. 




Chloride ~*v ^ 

Pot moistened 
with saturated 
solution of +. 

' acid 

Calcium Chloride £ _ _ 
Fig. 31. 

The apparatus usually supplied for this experiment (see Fig. 31) 
is far too large to be filled in a reasonable period of time, and is too 
liable to burst. A smaller model of about 50 cm 3 capacity (obtainable 
from manufacturers quite easily) is very suitable and the experiment 
can be completely performed in 20 minutes. 

Set up the apparatus of Fig. 31 in diffused daylight, and allow the 




mixed gases to pass for 5 to 10 minutes. Close both taps. Attach 
wires from a coil, and arrange a plate of thick glass between the tube 
and the observers. Close the circuit and a flash is seen but no sound 
is heard. Note that the green colour of chlorine has now disappeared. 
Fill the tube leading to one tap with mercury, and, holding the 
liquid in place with the finger, invert under mercury in a mortar. 
Open the tap. No gas enters or leaves. Close the tap, replace the 
mercury in the tube by water, and open under the surface of water. 
The latter rises and almost fills the tube. 

Note. The experiment on page 142 (Electrolysis of hydrochloric 
acid) shows that the gases which fill the eudiometer tube consist of 
equal volumes of hydrogen and chlorine. 

This proves that the whole of the gas in the tube was hydrogen 
chloride because any excess of hydrogen or chlorine would not have 
dissolved with this rapidity. This experiment proves that, starting 
with half a tubeful of hydrogen and a half a tubeful of chlorine 
(see p. 142), we obtain, by their combination, a tubeful of hydrogen 
chloride (temperature and pressure constant) or 

1 volume of hydrogen combines with 1 volume of chlorine to give 
2 volumes of hydrogen chloride (temperature and pressure 
Applying Avogadro's Law, we may substitute molecules for 
volumes, all the substances being gases. 

.-. 1 molecule of hydrogen combines with 1 molecule of chlorine 

to give 2 molecules of hydrogen chloride. 
But 1 molecule of hydrogen and 1 molecule of chlorine each con- 
tain 2 atoms. 

/. 2 molecules of hydrogen chloride contain 2 atoms of hydrogen 

and 2 atoms of chlorine 
.*. 1 molecule of hydrogen chloride contains 1 atom of hydrogen 
and 1 atom of chlorine 

/. The formula of hydrogen chloride is HC1 

The equation for the above reaction is H 2 + Cl a — > 2HC1. 
Note. Since the reactants and products are all gaseous the vapour 
density is not required. 

Ammonia gas 

The full argument depends upon two experiments (a) and (b). 

(a) Ammonia gas can be formed by sparking nitrogen and hydro- 
gen in suitable proportions. This proves that ammonia contains 
these two elements only. (This experiment is not described here.) 




(b) Hofmann's method (described below) may be used to demon- 
strate the volume proportions of the nitrogen and hydrogen combined 
in ammonia. The method depends on the following facts: 

(i) That ammonia reacts with chlorine, liberating nitrogen and 

forming hydrogen chloride, 
(ii) That the hydrogen and chlorine combine in equal volumes 

when hydrogen chloride is formed. (See last experiment.) 

Take an apparatus similar to that of Fig. 32. It can easily be made 
by heating a burette (a damaged one will be quite satisfactory if the 
length from the tap to the open end is 40 cm or more) and drawing 

out the heated portion until it forms a 

VC constriction of which the diameter is 

about that of ordinary glass-tubing, and 
then making a file-scratch on the con- 
striction and breaking AB separate from 
C. AB should be about 30 cm long 
and graduated into three equal portions. 
Attach AB to a chlorine generator 1 in a 
fume-chamber and fill it with chlorine. 
Add the rubber-tubing and clip, and the 
portion C, and clamp the whole appa- 
ratus in a vertical position. Put a little 
concentrated ammonia into C, cautiously 
release the clip and allow a little of the 
ammonia to enter AB. There will be a 
flash of light and white fumes of ammo- 
nium chloride will appear. Carefully allow 
more ammonia to enter AB until there 
is no further reaction and a few drops 
of liquid have collected at B. Then put dilute sulphuric acid into 
C, colour it with a little litmus solution and run it into AB to 
neutralise the excess ammonia, taking care that C does not become 
empty of liquid or air will be drawn into AB. When the solution in 
AB is red (i.e., when all the ammonia has been neutralised) remove 
C and place AB in a wide, deep vessel filled with water so that 
the tap B is well immersed. Open the tap. Water will enter. Push 
AB down into the vessel until the levels of water inside and out- 
side are equal, so giving atmospheric pressure inside AB. It will then 
be found that the gas left occupies one-third of the volume of AB. 
The gas is nitrogen. 

1 To teachers. It is convenient lo make this apparatus part of the delivery-tube 
when chlorine is being prepared as a demonstration experiment. The tube is 
then automatically filled with chlorine without extra trouble. 






Fig. 32. 
Formula of ammonia. 

The original volume, AB, of chlorine combined with hydrogen 
from the ammonia to form hydrogen chloride. A volume, AB, of 
hydrogen from the ammonia must have been used up to combine 
with this chlorine (see [ii] p. 1 14). At the same time, one-third of a 
volume AB of nitrogen was liberated from the ammonia. 

*. 3 volumes of hydrogen were combined with 1 volume of 
nitrogen in ammonia gas (temperature and pressure constant). 

Applying Avogadro's Law, we may replace 'volumes' by 'mole- 

.*. 3 molecules of hydrogen combine with 1 molecule of nitrogen 
.*, 6 atoms of hydrogen combine with 2 atoms of nitrogen or, 

in the simplest terms, 
3 atoms of hydrogen combine with 1 atom of nitrogen 
.*. the simplest (empirical) formula for ammonia gas is NH 3 and 
its molecular formula is (NH 3 )« where n is a whole number 

,'. the molecular weight of ammonia gas is (14 + 3)/i or 17n. 

We now need the molecular weight of ammonia gas to find the 
value of n. 

The vapour density of ammonia gas is 8.5. 

.', its molecular weight is 17 

/. 17n = 17 

/. n= 1 

.*. the molecular formula of ammonia gas is NH 3 


2NH 3 + 3C1 2 ->Nj-f-6HCl 
Then 6HC1 + 6NH 3 -> 6NH„Cl 

Adding 8NH 3 + 3C1, -*• N 2 + 6NH 4 C1. 


The apparatus (see Fig. 33) consists of a stout eudiometer tube 
surrounded by a jacket which contains a vapour at 130°C (amyl 
alcohol boils at 130"C at 760 mm pressure). The other limb of the 
eudiometer tube serves as a manometer, for mercury can be run 
into and out of this tube, so altering the pressure. One volume of 
oxygen is introduced and then two volumes of hydrogen. The open 
end of the tube is suitably plugged so that the mercury is not blown 
out. The mixture is exploded by means of an electrical spark and the 



plug is removed. On allowing the gas to cool down to 130°C and 
equalising the mercury levels in the two tubes it is found that there 
are two volumes of steam left in the tube. (All the above measure- 
ments of volume are made at laboratory pressure and 130°C.) On 
cooling the apparatus below 100°C the steam condenses to water 
and the mercury rises to the top of the enclosed tube showing that 
all the oxygen and hydrogen have been used up since they would not 
condense to a liquid as does the steam. 

From this experiment 

2 volumes of hydrogen and 1 

Vapour of 

Amy/ Alcohol 

7b Induction 



volume of oxygen form 2 vol- 
umes of steam. 

Applying Avogadro's Law we may 
substitute molecules for volumes, all 
the substances (at this temperature) 
being gaseous. 

.*. 2 molecules of hydrogen and 
1 molecule of oxygen form 2 
molecules of steam 

.'. 1 molecule of hydrogen and \ 
molecule of oxygen form 1 

Fl0, "• molecule of steam 

Volume composition of steam. 

But the molecule of hydrogen con- 
tains two atoms and the molecule of oxygen contains two atoms. 

.*. Formula for steam is H 2 

Note. Since the reactants and the products are all gaseous under the 
conditions of the experiment the vapour density is not required. 

Nitrogen monoxide 

A suitable volume of nitrogen monoxide is measured at atmos- 
pheric pressure (levels A and B equal) in the hard-glass tube (Fig. 34). 
The spiral of iron wire is then electrically heated to red heat. The 
metal combines with the oxygen of the nitrogen monoxide liberating 
nitrogen (the residual gas can be shown to be inactive). After about 
20 minutes the electrical current is switched off and the tube allowed 
to cool. Water-level A rises towards C. 

After the tube has been transferred to a deep vessel and lowered 
until the levels C and B are equal, the volume of nitrogen is measured 
and is found to be one-half of the original volume of nitrogen mon- 



That is, 1 volume of nitrogen is contained in 2 volumes of nitrogen 
monoxide (temperature and pressure constant). 

Using Avogadro's Law, we may substitute 'molecules' for 


Hard glass tubs 
Iron wire 


Fig. 34. 
Volume composition of oxides of nitrogen. 

.'. 1 molecule of nitrogen is contained in 2 molecules of nitrogen 

.'. 2 atoms of nitrogen are contained in 2 molecules of nitrogen 

,\ 1 atom of nitrogen is contained in 1 molecule of nitrogen 

/. the formula of nitrogen monoxide is NO* where x is a whole 

.'. the molecular weight of nitrogen monoxide is (14 -j- 16x). 

The vapour density of nitrogen monoxide is 15. 
The molecular weight of nitrogen monoxide is 30. 

.-. 14 + 16x = 30 
x= 1 
.". formula of nitrogen monoxide is NO 
3Fe + 4NO -»• Fe 8 4 + 2N, 



Note. Pure nitrogen monoxide for this experiment is conveniently 
prepared by half-filling a small flask with iron(II) sulplmte crystals, 
covering them with dilute sulphuric acid, warming and dropping 
sodium nitrite solution into the mixture from a tap-funnel. 

Dinitrogen oxide 

Exactly the same experiment is performed as for nitrogen monoxide 
except that the gas must be confined over mercury because dinitrogen 
oxide is fairly soluble in water. The hard-glass tube should be only 
about half-filled with the gas at first. It is found that the volume of 
nitrogen left is equal to the volume of dinitrogen oxide taken. 

.". 1 volume of nitrogen is contained in 1 volume of dinitrogen 
oxide (temperature and pressure constant). 

Using Avogadro's Law, we may substitute 'molecules' for 

.'. 1 molecule of nitrogen is contained in 1 molecule of dinitrogen 

.". 2 atoms of nitrogen are contained in 1 molecule of dinitrogen 

.". the formula of dinitrogen oxide is N s O a (* is a whole number) 
.'. the molecular weight of dinitrogen oxide is (28 + 16x). 
The vapour density of dinitrogen oxide is 22 
,*. Molecular weight of dinitrogen oxide is 44. 
.*. 28 + \6x = 44 

:.x= l 

.*. the formula of dinitrogen oxide is N 2 
3Fe + 4N 2 -*■ Fe 3 4 + 4N„ 

Hydrogen sulphide 

A convenient volume of hydrogen sulphide is confined over mer- 
cury at atmospheric pressure (Fig. 35). By means of an induction 
coil, electric sparks are passed between the platinum wires for some 
time. This decomposes the hydrogen sulphide into its elements. 
Solid sulphur (of negligible volume) is deposited and hydrogen is 
left. It is found that, when the tube has cooled, the volume of hydro- 
gen left is exactly equal to the volume of hydrogen sulphide taken 
(temperature and pressure constant). 

.". 1 volume of hydrogen is contained in 1 volume of hydrogen 
sulphide (temperature and pressure constant). 








Fia. 35. 

Volume composi- 
tion of hydrogen 

.". the 

Applying Avogadro's Law, we may substitute 
'molecules' for 'volumes'. 

.". 1 molecule of hydrogen is contained in 1 
molecule of hydrogen sulphide 

.". 2 atoms of hydrogen are contained in 1 
molecule of hydrogen sulphide 

.". the formula of hydrogen sulphide is rlgS*, 
where at is a whole number 

,*. the molecular weight of the gas is (2 + 32x). 

The vapour density of hydrogen sulphide is 17. 

.'. The molecular weight of hydrogen sulphide 
is 34. 

/. 2 + 32x = 34 
x= 1 
formula of hydrogen sulphide is H g S 
H 2 S-»-H 2 + S 
Carbon dioxide 

The apparatus of Fig. 36 is used. A little dry powdered graphite 
(0.02 g is required for 40 cm 3 ) is placed in the dry tube, and dry 
oxygen is passed through the whole apparatus for 3-4 minutes. The 
clip is closed and mercury poured into the manometer. The clip is 
momentarily released to make the pressure in the tube atmospheric. 
The graphite is heated with a small flame, and it burns (not always 
obviously) to carbon dioxide. Expansion causes the level at C to fall, 
and the level at D to rise, 1 but, when the bulk has cooled, the levels 
at C and D return to their original positions, that is, the volume of 
carbon dioxide formed is equal to the volume of oxygen used. 

.'. 1 volume of oxygen is contained in 1 volume of carbon dioxide 

(temperature and pressure constant). 
Applying Avogadro's Law, we may substitute 'molecules' for 

.". 1 molecule of oxygen is contained in 1 molecule of carbon 


.*. 2 atoms of oxygen are contained in 1 molecule of carbon dioxide 

.'. its formula is C„O s 

.'. its molecular weight is 12* + 32. 

* If the level at C falls so that gas is in danger of being lost, remove the 
burner and close the mouth of tube D lightly with the finger. 


The vapour density of carbon dioxide is 22. 
.". The molecular weight of carbon dioxide is 44. 
/. 12* + 32 = 44 
x = 1 
.'. the formula of carbon dioxide is CO a 
C + o 2 -> C0 3 


Dry Carbon 





Fig. 36. 
Formula of carbon dioxide. 

Sulphur dioxide 

The same experimental work is carried out as for carbon dioxide, 
substituting sulphur for carbon. As in the case of the carbon dioxide 
the volume of sulphur dioxide formed is equal to the volume of 
oxygen used. 

The reasoning to obtain the formula, S0 2 , is the same as for 
carbon dioxide, using the vapour density of sulphur dioxide, 32, 
and the atomic weight of sulphur, S = 32. 


Ozone can be made from oxygen and can be converted by heat 
into oxygen and no other product. It must therefore have the formula, 
O n . The value of n is established as follows: 

1. Ozone reacts with turpentine. Suppose on treatment with tur- 
pentine, 200 cm 3 of ozonised oxygen shrink in volume by x cm 3 . 
This means that in the 200 cm 3 of ozonised oxygen there are x cm 3 of 



2. Another 200 cm 3 of the same sample of ozonised oxygen is 
heated, converting the ozone to oxygen. It will be found that there is 
an increase in volume of x/2 cm 3 on cooling to the original tempera- 
ture, pressure remaining constant. This means that the oxygen formed 
from the ozone occupies the x cm 3 formerly occupied by the ozone, 

together with a further x/2 cm 3 , i.e., — cm 3 in all. 

.-. x cm 3 of ozone yield — cm 3 oxygen at same temperature and 

Or 2 volumes of ozone yield 3 volumes oxygen at same temperature 
and pressure. 
.*. By Avogadro's Law, 

2 molecules of ozone yield 3 molecules of oxygen, 
i.e., 20„ = 30, 

.-. n = 3 
and the formula of ozone is O s 

Carbon monoxide 

A measured volume of carbon monoxide, confined over mercury in a 
eudiometer tube, is mixed with a measured volume of oxygen equal to 
several times its own volume. The mixture is exploded by a spark passed 
between platinum leads sealed through the glass. After cooling, the volume 
of residual gas is measured and the carbon dioxide is absorbed by allowing 
some concentrated caustic potash solution to rise above the mercury. (The 
potash solution is introduced at the bottom of the tube by means of a small 
pipette bent at the tip.) The diminution in volume caused by the potash 
represents the volume of carbon dioxide formed. The residual volume is 
the excess oxygen, and, by subtracting this from the original volume of 
oxygen taken, the volume of oxygen used up is obtained. (All measurements 
are taken at room temperature and atmospheric pressure.) 

It will be found that : 
2 volumes of carbon monoxide combine with 1 volume of oxygen to form 2 
volumes of carbon dioxide. 

Using Avogadro's Law, we may substitute 'molecules' for 'volumes'. 
Then, 2 molecules of carbon monoxide combine with 1 molecule of oxygen 
to form 2 molecules of carbon dioxide. 

.". 1 molecule of carbon monoxide contains \ molecule less oxygen than 

one molecule of carbon dioxide 
.*. 1 molecule of carbon monoxide contains one atom of oxygen less 

than one molecule of carbon dioxide 
But the formula of carbon dioxide is COj 

.'. the formula for carbon monoxide is CO. 
Note. Since the reaclants and products are all gaseous the vapour density 
is not required. 




1. How would you show experimentally that two volumes of hydrogen 
combine with one volume of oxygen to yield two volumes of steam? What 
deductions can be drawn from these facts as to the number of atoms in the 
oxygen molecule? (C.) 

2. Complete the following statements: 

(a) Two volumes of hydrogen unite with volume(s) of oxygen to 

give volume(s) of steam, all the substances being at 100°C and 

atmosphere pressure. 

(b) Two volumes of hydrogen unite with volume(s) of nitrogen to 

give volumes) of ammonia, all the gases being at the same 

temperature and pressure. 

State the law which these statements illustrate, and give the name of its 

Explain clearly how the formula; of steam and ammonia follow from 
the statements you have made. (N.U.J.B.) 

3. Describe two experiments from the results of which the formula CO, 
for carbon dioxide may be derived. Point out clearly what assumptions are 
made in the deduction of the formula. (B.) 

4. Describe a eudiometer and explain shortly how it is used. Give an 
account of the experiments by which the formula of ammonia may be 
shown to be NH,. (L.) 

5. Describe the preparation and principal properties of carbon mon- 
oxide. What is the evidence on which the formula CO is assigned to this 
gas? (L.) 

6. Give an account of the experimental evidence on which the accepted 
formula for hydrogen chloride is based. 223 cm 3 of hydrogen chloride and 
250 cm" of gaseous ammonia, both measured at 13°C and 770 mm, were 
mixed. Calculate the weight of the solid product. (CI - 35.5, N = 14.) 

7. State Gay-Lussac's Law of Gaseous Combination, illustrating your 
answer by reference to four examples. 

How would you show experimentally that a given volume of hydrogen 
sulphide contains twice as much hydrogen as an equal volume of hydrogen 
chloride? (L.) 

Chapter 1 1 

Volumetric Analysis 

The volume units employed for solutions in this section 
are the cubic decimetre (dm 3 ) and the cubic centimetre (cm 3 ). 
They are respectively equivalent to the litre (1) and millililre 
(ml), the use of which is discouraged in the S.I. system. 

VOLUMETRIC analysis is a means of estimating quantities of 
certain materials (often acids or alkalis) by an analytical process 
which involves measurement of volumes of solutions, using pipettes, 
burettes and (for approximate measurements) measuring cylinders. 
Weighings may also be involved. 

Definition. A standard solution is a solution of which the concentration 
is known. 

For example, a solution known to contain, say, 12 g of sodium 
chloride in one dm 3 of solution is a standard solution. 

The system now generally approved for volumetric work is based 
upon the molar (M) solution. Such a solution of a given compound 
contains the molecular weight in grams (or one mole) of the com- 
pound in one cubic decimetre (dm 3 ) of the solution. 
Definition. A molar (M) solution of a compound, X, is one which 
contains one mole of the compound in one cubic decimetre (dm 3 ) of 
the solution. 

For example, since QHigO,, = 180, a molar solution of glucose 
contains 180 g in 1 dm 3 . In the case of strong electrolytes such as 
sodium chloride or sodium carbonate, which do not exist as mole- 
cules to any significant extent in dilute solution, the expression 
gram formula weight is often used as equivalent to one mole. For 
example, the gram formula weight of NajjCOg is taken as 106 g 
(Na 2 C0 3 = 106) and molar sodium carbonate solution contains 
106 g of the anhydrous salt in one dm 3 of solution. In fact, the 
dissolved salt is almost entirely dissociated into ions (in dilute 
solution) in the proportion of 2Na + to C0 3 2_ . Also, with many 
salts, hydration of ions occurs in solution. 




Molar solutions of some compounds commonly used in titration 
contain the following weights of the compounds in 1 dm 3 of solution: 

sodium hydroxide, NaOH 40 g 

potassium hydroxide, KOH 56 g 

sulphuric acid, H 8 S0 4 98 g 

hydrochloric acid, HC1 36.5 g 

sodium carbonate, Nas.CO a 106 g 

sodium hydrogen carbonate, NaHCO, 84 g 

The figures quoted arc the gram-molecular weights (gram formula 
weights) of the respective compounds. 

Derivative concentrations are also used, e.g., 0.1 M, 0.5 M, 2 M, 
and these contain one-tenth of, half of and twice (respectively), the 
weight of solute present in an equal volume of the corresponding M 

It is obvious that some accurately standard solution is required as 
a starting-point for volumetric estimations. Actually few compounds 
are suitable for the direct preparation of an accurately standard 
solution. Some compounds absorb water from the air and, for this 
reason, cannot be weighed out accurately without excessively 
difficult precautions, e.g., sodium hydroxide, potassium hydroxide 
and concentrated sulphuric acid. Others react with carbon dioxide 
of the air, e.g., 

2NaOH (or 2K.OH) + CO a -► Na 8 CO s (or K 2 CO s ) + H.O 
Solutions may contain volatile constituents and be liable to change 
slowly in concentration during ordinary use, e.g., concentrated 
hydrochloric acid and ammonia. 

A compound which is commonly utilised for the direct preparation 
of an accurately standard solution is anhydrous sodium carbonate. 
It is best made from sodium hydrogen carbonate of high purity 
(which can be bought from chemical suppliers). This is done by 
heating the sodium hydrogen carbonate to constant weight, which 
ensures completion of the decomposition: 

2NaHCO a -»- Na,CO a + H t O + CO s 
The anhydrous sodium carbonate so formed is very pure and can 
be used in ordinary weighings with no appreciable change of com- 

Preparation of exactly one-tenth molar (0.1 M) sodium carbonate 

As stated above, molar sodium carbonate solution contains 106 g 
of solute in 1 dm 3 (since Na a CO s = 106). Consequently, a 0.1 M 
sodium carbonate solution contains 10.6 g of solute in 1 dm 3 . The 



usual volume of solution made up in a laboratory experiment is 
250 cm 3 because this volume allows of three titrations of 25 cm 3 
each, with a reserve for required washings of pipette and possible 
titration failures. For 250 cm 3 of 0.1 M solution, the weight of an- 
hydrous sodium carbonate required is one-quarter of 10.6 g, or 

2.65 g. 
Experimental. Put about 8 g of pure sodium hydrogen carbonate 

into a dry, clean evaporating dish and heat it with stirring for, say, 

15 minutes. Place it in a desiccator to cool, then weigh it. Without 

further stirring (to prevent loss of solid), heat the dish again for 

five minutes or so, cool it in the desiccator and weigh, and so on 

until two consecutive weighings are the same. Store the dish and pure 

sodium carbonate in the desiccator. 

Weigh a watch-glass (or, better, a stoppered weighing bottle) and 
weigh out, on to it, exactly 2.65 g of anhydrous sodium carbonate. 
Tap the carbonate into a beaker (with lip) of about 400 cm 3 capacity, 
containing about 50 cm 3 of hot distilled water. Wash down the 
watch-glass (or weighing bottle) with a jet of hot distilled water 
from a wash-bottle and allow the washings to fall into the beaker. 
This procedure should wash all the sodium carbonate into the 
beaker. Stir with a glass rod till dissolution of the solid is complete, 
then cool the solution to room temperature. Leave the rod standing 
in the solution. 

Place a very thin smear of Vaseline under the lip of the beaker 
and pour the solution down the glass rod into a measuring flask 
(250 cm 3 ). Wash the beaker out at least twice with jets of cold dis- 
tilled water directed round the sides and pour the washings down 
the glass rod into the measuring flask. Shake the flask gently. Fill it 
up with cold distilled water almost to the mark, then add more dis- 
tilled water drop by drop from a pipette till the lowest level of the 
meniscus is on the mark when at eye-level (to avoid parallax error). 
Stopper the measuring flask and shake well. The liquid should then 
be exactly 0.1 M sodium carbonate solution. 

Preparation of a 0.1 molar solution of sulphuric acid 

This preparation cannot be done directly because concentrated 
sulphuric acid absorbs water rapidly from the air and is never 
reliably pure. A solution is prepared which is a little above 0.1 M in 
concentration and it is then standardised and diluted with distilled 
water to exactly 0.1 M. 

A molar solution of sulphuric acid contains 98 g of pure acid in 
1 dm 3 , so the 0.1 M acid contains 9.8 g of the acid in 1 dm 3 . The 
concentrated acid has a density of about 1.8 g/cm 3 , so 9.8 g of it 
occupy about 5.5 cm 3 . 



25 cm 3 
mark U 


Experimental. With great care, because of the dangerously cor- 
rosive nature of the acid, take 5.5-6.0 cm 3 of concentrated sulphuric 
acid in a small measuring cylinder; pour it, with stirring, into about 
100 cm 3 of cold distilled water in a beaker (with lip). Pour this 
solution into, say, 700 cm 3 of cold distilled water in a measuring 

flask of capacity 1000 cm 8 , 

wash out the beaker with 

cold distilled water twice 

1 — » — and add the washings to the 

measuring flask. Then add 
distilled water approximately 
to the mark on the measuring 
flask, stopper it and shake 
well. This should give sul- 
phuric acid of concentration 
a little above 0.1 M. It is 
now standardised with the 
0. 1 M sodium carbonate sol- 
ution prepared above. 

Wash out a burette 
(50 cm 3 ) twice with a few cm 8 

50 cm » £. of the approximately 0.1 M 

sulphuric acid and run the 
acid out through the tap. 
(This leaves the burette wet 
with the liquid it is to contain 
and nothing is left to conta- 
minate the acid solution when 
the burette is filled.) Fill the 
burette 1 cm or so above the 
cm 3 mark and run a little 
of the acid out through the 
tap, so bringing the acid level 
a little below the cm 3 mark 
and filling the tip of the 
burette with acid. 


Fio. 37. 
Burette and pipette. 

Wash out a pipette (25 cm 3 ) twice with a little of the exactly 
0.1 M sodium carbonate solution. This leaves it wet with the solution 
it is to measure. Draw this solution into the pipette above the mark 
and, with the mark at eye-level, allow the solution to run out very 
slowly till the lowest level of the meniscus is on the mark. Then allow 
the liquid to run from the pipette into a conical flask (Fig. 37) which 
has been washed out with cold distilled water only and touch the tip 
of the pipette on to the surface of the liquid. (Do not blow out the 



last drop.) The flask will then contain exactly 25 cm 3 of 0.1 M sodium 
carbonate solution (with a little distilled water which is of no 

Add two drops of methyl orange solution (indicator) to the conical 
flask. This will turn the alkaline solution yellow. Read and note the 
level of acid in the burette. Run acid into the conical flask 2 or 3 
cm 8 at a time, with shaking, allowing a short drainage time between 
additions, until the liquid flashes pink. Then add the acid more slowly, 
eventually drop by drop, until the colour of the liquid is orange {i.e., 
the neutral colour for methyl orange). Slight excess of acid will turn 
it fully pink. Read and note the acid level. Repeat the titration to 
obtain at least two results which differ by not more than 0.1 cm'. 


25.0 cm 8 of 0.1 molar sodium carbonate solution used each time 

Titration 1 2 3 

2nd burette reading 24.0 23.9 24.0 cm 3 

1st burette reading 0.2 0.3 0.4 cm 8 

Acid added 218 216 216 cm 8 

The first titration can be neglected as a trial run and the average of 
the other two is 23.6 cm 8 . 

Since 25.0 cm 8 of the exactly 0.1 M sodium carbonate solu- 
tion require only 23.6 cm 3 of the acid, the acid must be the more 
concentrated in the proportion of 25.0 : 23.6, i.e., the acid must 

be — - x 0.1 M or 0.106 M. The concentration of acid is 98 X 

0.106 g/dm 3 , or 10.4 g/dm 8 . To make the acid exactly equal in molar 
concentration to the sodium carbonate solution, 23.6 cm 8 of the acid 
must be diluted to 25.0 cm 3 , i.e., 1.4 cm 3 of distilled water must be 
added to 23.6 cm 8 of the acid. If 920 cm 3 of the acid are left, it can 

1 4 
be made exactly 0.1 M by the addition of 920 X -^ cm 8 of distilled 

water, or 55 cm 8 . After the addition of this volume of distilled water, 
the acid should be shaken and should then be exactly 0.1 M. This 
can be tested by titration of the 0.1 M sodium carbonate solution, 
when 25.0 cm 3 of the acid should be required to neutralise 25.0 dm 8 
of the carbonate solution. 

These two standard alkaline and acidic solutions can be used 
to standardise other solutions, e.g., sodium hydroxide, potassium 
hydroxide, hydrochloric acid, and, by their use, a wide range of 
acid-alkali estimations can be carried out. The following examples 
illustrate these estimations. 



27.5 cm 3 of a solution of sodium hydroxide neutralise 25.0 cm 3 of 
M hydrochloric acid. Calculate (a) the molarity, (b) the concentra- 
tion of the sodium hydroxide solution in g/dm 3 . 

The object of the calculation is to fix the molarity and concentra- 
tion of the sodium hydroxide solution. We know the molarity of the 
hydrochloric acid; consequently we require to know neither the 
molecular weight of the acid nor its concentration in g/dm 8 . 
NaOH + HC1 -^NaCl + HgO 
ldm 3 ofM ldm 3 ofM 
The equation shows that M sodium hydroxide solution and M hydro- 
chloric acid are equivalent to one another volume for volume. Because 
it requires the smaller volume in the titration, the acid used is the 
more concentrated (in molar terms) of the two reagents. That is, the 

sodium hydroxide solution is — ^- M, or 0.909 M. 

Since the molecular weight of sodium hydroxide is 40, the solution 
must contain 0.909 X 40 g/dm s , or 36.4 g/dm 3 . 

50 cm 3 of M sulphuric acid are added to an excess of solid sodium 
hydrogen carbonate. Calculate (a) the weight of sodium sulpliate pro- 
duced, (b) the volume of carbon dioxide evolved at \5°C and 770 mm 
pressure. {NOpSOf = 142; the molecular weight in grams of any gas 
occupies 22.4 dm 3 at s.t.p.) 

This calculation requires the equation for the reaction and the 
insertion into it of the quantities of the reagents involved in the units 
in which they are expressed in the data given. The sodium hydrogen 
carbonate is only vaguely stated as in excess and can be ignored 
quantitatively. The acid is stated in terms of molarity and is so ex- 
pressed in the equation. Sodium sulphate is required in grams and 
the carbon dioxide in terms of volume. The required reaction state- 
ment then becomes: 

2NaHCO s + H,S0 4 -> Na,S0 4 + 2H 2 + 2CO, 

1 dm 3 of M 142 g 2 x 22.4 dm 3 at 


From the equation, 

1000 cm 3 of M HjSO, produces 142 g of sodium sulphate 

so 50 cm 3 of M H 2 S0 4 produces 142 x — — g of sodium sulphate, 


i.e., 7.1 g 
From the equation, 

1000 cm 3 of M H t S0 4 produces 44.8 dm 3 of CO, at s.t.p. 



so 50 cm 3 of M H 2 S0 4 produces 44.8 X — - dm 3 of CO,, or 


2.24 dm 3 at s.t.p. 
Converted to 15°C and 770 mm pressure, this volume becomes 


or 2.33 dm 3 . 

15.0 g of anhydrous sodium carbonate, containing some sodium chlor- 
ide as impurity, were made up to 250 cm 3 of solution and 25.0 cm 3 of 
the solution were titrated by M HCI. 24.5 cm 3 ofM HCl were required 
to neutralise the liquid. Calculate the percentage of sodium chloride in 
the solid. 

For this calculation, the equation is required, with the insertion of 
the quantity of hydrochloric acid in terms of molarity and of sodium 
carbonate in grams. Sodium chloride does not react. 

Na,CO s + 2HC1 -> 2NaCl + H,0 + CO a 
106 g 2 dm 3 of M 

Wt. of sodium carbonate in 25.0 cm 3 of solution = 

106 X 




Wt. of sodium carbonate in 250 cm 3 of solution = 

Therefore, wt. of sodium chloride in the mixture is 
(15.0 -1 3.0) g, or 2.0 g 

That is, percentage of sodium chloride = -r^T. x 10 ° 

= 13.3 

Normal Solutions 

The normality system of concentrations for volumetric analysis, 
which has been in use for many years, is based on the normal (N) 

Definition. A normal (N) solution of a compound, X, is one which 
contains the equivalent weight in grams of X in one litre of solution. 

The terms equivalent and litre have recently been declared scienti- 
fically obsolete after international discussions but the normality sys- 
tem may continue in use for some time yet. 

Acids. The gram-equivalent weight of an acid is that weight of the 
acid in grams which produces 1.008 grams of hydrogen ion, H + . 



From the following ionisations: 

HC1 ^= H+ +C1- 

36.5 g 1.008 g 

H 2 S0 4 ^ 2H+ + SO« 8 - 
98 g 2x 1.008 g 

it is obvious that the gram-equivalent weight of hydrogen chloride is 
36.5 g, and of sulphuric acid is 98/2 g, or 49 g. That is, a normal 
solution of hydrochloric acid contains 36.5 g, and of sulphuric acid 
49 g, of acid in one litre. 

Alkalis and carbonates. The gram-equivalent weight of an alkali (or 
carbonate) is that weight in grams of the alkali (or carbonate) that 
reacts with 1.008 g of hydrogen ion. 

From the following reactions : 

NaOH (or KOH) + H + 
40 g 56 g 1.008 g 

Na^COj + 2H + 

106 g 2 X 1.008 g 

Na + (or K + ) + H a O 
2Na + + H t O + C0 8 

it is obvious that the gram-equivalent weight of sodium hydroxide is 
40 g, of potassium hydroxide is 56 g and of Na 2 C0 3 is 106/2 g, or 
53 g. That is, a normal solution of sodium hydroxide contains 40 g 
of the alkali in one litre; for potassium hydroxide, the corresponding 
weight is 56 g and, for sodium carbonate, it is 53 g. 

A 0.1 N solution of sodium carbonate can be made up accurately 
as described for the 0.1 M solution earlier in this chapter. The weight 
of sodium carbonate needed for 250 ml of 0.1 N solution is 53 X 0.1 
X 250/1000 g, or 1.325 g. Similarly, an approximately 0.1 N solution 
of sulphuric acid can be made up as described earlier and requires 
2.8-3.0 ml of concentrated sulphuric acid in one litre. The acid can 
then be standardised as described for the 0.1 M acid. 

The normality system possesses the very useful characteristic that 
equal volumes of all normal solutions are chemically equivalent to 
one another. For example, 25.0 ml of N sodium hydroxide solution, 
or N sodium carbonate solution, are exactly neutralised by 25.0 ml 
of any strong acid. This situation does not apply to molar solutions. 
For example 25.0 cm s of M sodium hydroxide solution require 
25.0 cm 3 of M HC1 but only 12.5 cm 8 of M H,SO«. The monobasic 
and dibasic character of the acids (respectively) accounts for this 

The following examples illustrate calculations with normal solu- 

25.0 ml of a normal sodium carbonate solution are neutralised by 
24.5 ml of dilute hydrochloric acid. Calculate the normality of the acid 



and its concentration in grams /lit re. If 900 ml of the acid remain, how 
may it be made exactly normal? 

Since it needs a smaller volume, the acid is the more concentrated 
(in terms of normality) in the ratio of 25.0 : 24.5, i.e., the acid is 
25.0/24.5 N, or 1.02 N. Substituting the gram equivalent weight of 
HC1 (36.5 g) for N, the concentration of the acid in grams/litre is 
1.02 X 36.5 or 37.2. 

To make the acid exactly normal, it must become equivalent, 

volume for volume, to the sodium carbonate solution; that is, 0.5 ml 

of distilled water must be added for each 24.5 ml of acid. For 900 ml 

of acid, the volume of distilled water required is 0.5 X — ml, or 

18 ml. After addition of this water, the mixture should be shaken. 

Since the volume (r) of a solution required for a given purpose 
decreases in the same ratio as the normality (N) of the solution in- 
creases, it follows that the product, If X v, is constant. That is, if v , ml 
of a solution of normality, N x , are neutralised by v 2 ml of a solution 
of normality, N t , 

This relation is used in the example below. 

14.0 g of a sample of anhydrous sodium carbonate {containing some 
sodium sulphate as impurity) was made up to 250 ml of solution. 25.0 ml 
of this solution required 24.8 ml ofN sulphuric acid for titration. Cal- 
culate the percentage of impurity present in the carbonate. 

The sodium sulphate impurity does not react with the sulphuric 
acid. Applying the relation mentioned above, 

25.0 X iV, = 24.8 X N (where N = 1) 
the normality of the sodium carbonate solution being N t while the 
acid is normal. From this, 

^ = |^=0.992 

That is the alkaline solution is 0.992 N in sodium carbonate. Sub- 
stituting the gram equivalent weight of sodium carbonate (53) for N, 
the solution contains 0.992 X 53 g/litre, or 52.6 g. 

Consequently, the weight of sodium carbonate in 250 ml of solu- 
tion is 52.6/4 g, or 13.15 g. From this, 

percentage purity of sodium carbonate = -j^— X 100 

= 93.7 

so, percentage of impurity = 6.3 




1. 12.5 cm' of 0.5 M sulphuric acid neutralise 50.0 cm' of a given solu- 
tion of sodium hydroxide. What is the molar concentration of the sodium 
hydroxide solution? 

2. 25.0 cm' of a solution of sulphuric acid required 32.0 cm' of 0.1 M 
sodium hydroxide solution for neutralisation. Calculate the concentration 
of the acid in g/dm». 

3. 25.0 cm' of 0.05 M sulphuric acid neutralised 35.0 cm' of a potassium 
hydroxide solution. What is the concentration of the alkali in terms of 
(a) molar concentration, (b) g/dm'? 

4. 25.0 cm' of 0.5 M sulphuric acid arc mixed with 30.0 cm' of M 
sodium hydroxide solution. What volume of 0.1 M sulphuric acid will just 
neutralise the excess of alkali? 

5. What volume of a molar solution of sulphuric acid would exactly 
neutralise 25.0 cm* of a sodium hydroxide solution containing 60 g of 
the alkali in 1000 cm'? (Na = 23, O = 16, H = 1.) 

6. What volume of 0.1 M hydrochloric acid would react exactly with 
25.0 cm* of a sodium carbonate solution containing 5.20 g of anhydrous 
salt in 1 dm' of solution? (Na,CO, = 106.) 

7. 25.0 cm* of a solution containing 10.6 g of sodium carbonate in 
1000 cm* required 23.0 cm* of hydrochloric acid for neutralisation. 
Calculate the molar concentration of the acid and its concentration in 
g/dm*. (Na,CO, - 106; HC1 - 36.5.) 

8. 25.0 cm* of a solution of sodium hydroxide containing 10.0 g in 
1 dm* required 24.0 cm' of dilute sulphuric acid for neutralisation. Cal- 
culate the molar concentration of the acid and its concentration in g/dm*. 
(NaOH = 40, H,S0 4 = 98.) 

9. 1.00 g of pure ammonium chloride was boiled with 20.0 cm' of a 
solution of sodium hydroxide until the evolution of ammonia had ceased. 
If the resulting solution required 1 1.0 cm* of 0.1 M hydrochloric acid for 
neutralisation, calculate (a) the volume of ammonia gas evolved at 
s.t.p., (b) the molar concentration of the sodium hydroxide solution. 
(NH 4 C1 = 53.5; one mole of any gaseous compound occupies 22.4 dm* 
at s.t.p.) 

10. Describe how you would transfer exactly 25 cm* of a molar solution 
of sodium carbonate to a conical flask, using a pipette. How would you 
then set up and fill a wet burette with a given sample of dilute sulphuric 
acid and titrate the acid against the sodium carbonate solution? Mention 
throughout the precautions you would observe to secure accuracy and state 
the name and colour change of the indicator you would use. If 37.5 cm' 
of acid were needed for the titration, what is the molar concentration of 
the acid? 

11. Calculate how many cm' of 0.5 M sulphuric acid will be used in the 
titration of 25.0 cm' of a solution which contains 70 g of sodium hydrogen 
carbonate in 1 dm' of solution. State the molar concentration of the sodium 
hydrogen carbonate solution. (H — 1, Na — 23, C — 12, O — 16.) 

12. A typical sample of concentrated hydrochloric acid has a density 
of 1.16 g cm - ' and contains 32.0% of hydrogen chloride. What volume of 
this liquid would be needed to make 2 dm* of M hydrochloric add? 

Chapter 12 

Electrochemical Series and Electrolysis 

+ Cu 
+ 2Ag 
+ 2Ag 


Method of obtaining the series 

IT is a well-known fact that certain metals will displace other metals 
from solutions of their salts in water. For example, iron will displace 
copper from copper sulphate solution, and zinc will displace silver 
from silver nitrate solution. 

Fe -f- CuS0 4 -> FeS0 4 

Fe + Cu* + -*■ Fe 2+ 

Zn + 2AgN0 3 -*■ Zn(N0 3 ) a 

Zn + 2Ag+ -v Zn 2+ 
We can arrange the metals in a series such that any metal higher up 
in the series will displace from its salts any metal below it. The greater 
the gap separating the metals in the series, the more readily does 
displacement take place. 

Again, if a plate of zinc and a plate of copper are immersed in dilute 
sulphuric acid, a current will flow from the copper to the zinc outside 
the cell and from the zinc to the copper inside the cell, if the plates 
are connected by a wire. Thus the copper and the zinc must be at 
different potentials when in contact with dilute sulphuric acid. These 
potentials can be measured and by arranging the metals according to 
this potential difference the same series is obtained as by the displace- 
ment method. The list obtained is as follows, omitting the less 
common metals: 


Potassium K Most electropositive metal. 

Calcium Ca 

Sodium Na 

Magnesium Mg 

Aluminium Al 


















Au I 

Least electropositive metal. 

Hydrogen, although not a metal, is placed in the series to indicate the 
position it would occupy. 

It will probably strike you at once that the metals occurring above 
hydrogen liberate that element from acids with an ease indicated by 
the interval separating the metal from hydrogen in the series. Thus 
magnesium and zinc liberate hydrogen readily (so would sodium and 
potassium, with such an 'ease' that the experiment would be danger- 
ous), whereas copper, which is below hydrogen in the series, does not 
liberate hydrogen from acids at all. 

Chemical activity of the more electropositive elements 

Metals such as sodium and potassium, which occupy positions high 
up in the series, are said to be very electropositive. These metals are 
very active chemically, and metals lower in the series are less active. 
In modern terms, this means that the more electropositive metals 
ionise readily by loss of electrons, e.g., 

Na-e--)-Na + ; Ca - 2e~ -* Ca ,+ 
The less electropositive metals ionise much less readily. One result of 
this relation is that, if conditions are similar and neither metal is 
affected by water, a more electropositive metal will displace a less 
electropositive metal from its salt in solution, e.g., 

Zn + CuSO« —*■ ZnSO« + Cu 
The essential feature of this relation is that the zinc atom (more 
electropositive) transfers two electrons to the copper ion in solution, 
which is converted to a copper atom and precipitated. The changes 

Zn -»■ Zn*+ + 2e~; Cu s+ + 2e~ -> Cu 
or, added together, 

Zn + Cu 2+ -> Zn*+ + Cu 

The zinc ion is left in solution in association with the SO«* - ion, as 
the salt, zinc sulphate. For galvanic couple, see p. 151. 




- e~ 

— »■ 

Na + ; 




Zn 2+ 



Al 3 '; 

The behaviour of metals in liberating hydrogen from water or 
dilute acid is a special case of this relation. All metals which are more 
electropositive than hydrogen displace it; the most electropositive 
metals, e.g., Na, K, Ca, displace hydrogen from water. As electro- 
positive nature decreases, metals require dilute acid, e.g., Zn, Fe, 
while lead (very close to hydrogen in the electrochemical series) 
requires hot, concentrated hydrochloric acid. In all these cases, the 
metallic atom ionises by supplying a number of electrons equal to its 
valency to hydrogen ions present in water or acid. These ions are 
converted to hydrogen atoms and, by pairing, to molecules which are 
liberated. For example, 

H + + e~ -> |H 2 (water) 
2H + + 2e~ -+■ H 2 (dilute HC1) 
3H+ + 3e- ->- liH, (dilute HC1) 
Copper and silver are less electropositive than hydrogen. They cannot 
displace it from water or acid. They are attacked by no mineral 
acid except a strongly oxidising compound such as nitric acid or hot, 
concentrated sulphuric acid. 

In general the more electropositive metals oxidise readily, while 
the least electropositive tend to be inactive; for example, copper and 
silver are not readily attacked by the oxygen of the atmosphere. They 
have been found as free metals in the earth's crust and used as coinage 
metals. The more electropositive metals do not occur free in nature 
but only as compounds, such as sodium chloride, zinc sulphide and 
aluminium oxide. 

Gradation in properties of compounds of metals according to position 

in the series 

Not only does the series give us a good estimate of the chemical 
activity of the metal itself, but in many cases the properties of com- 
pounds of the metals are graded according to the position of the metal 
in the series. Thus the nitrates of sodium and potassium on heating 
decompose to the nitrite, the remainder of the metals of the series as 
far as copper (inclusive) form nitrates which decompose into the 
oxides on heating, liberating nitrogen dioxide and oxygen. The lowest 
members of the series, mercury, silver and gold, form nitrates which 
decompose to the metal on heating. (See p. 518.) 

The following table shows some of these properties of the metals 
and their compounds. It will be noticed that, as a rule, a change 
occurs in the region of calcium and another change in the region of 
copper. Be careful not to be too dogmatic about the properties as 
indicated in this way. Some metals show 'false' behaviour by ac- 
quiring, in air, a coherent coating of oxide which renders them 



a 1 



I S3 









t/j —j 

if flip 




si * 

Z§ ° 


z as 












i «« 

ra.o — 5 
s -o so 


4 1 

■O 60 

£ J « 

< s 


9 g.° a S 

o s — ffl_- 







* <5 z" 

f 5 < 



oe 3 






inactive, e.g., ordinary aluminium foil is inert towards air, water or 
dilute sulphuric acid. If, however, it is dipped into mercuric chloride 
solution, a layer of amalgam forms and an oxide layer cannot cohere 
on the foil. Then the metal attacks both water and acid, and rapidly 
throws off flakes of oxide in air at ordinary temperature. 

2A1 + 6H + ->- 2A1 S+ + 3H,; 4A1 + 30, -► 2AI,0 3 

The place of non-metals in the series 

A list has been worked out to include common non-metals and is 
given below. 





In this complete series, the 
further apart two elements arc, 
the more likely they are to 
form a stable compound. Thus 
oxygen combines very readily 
with sodium and potassium. 
Elements close to one another 
either do not combine at all or 
form unstable compounds. 
Thus chlorine dioxide, C10„ 
is an unstable explosive sub- 
stance. Metals do not form 
stable compounds with each 










Electrolysis is the decomposition of a compound (molten or in 
solution) by the passage through it of an electric current. 

Electrolytes and non-electrolytes 

Consider the apparatus of Fig. 38. By varying the liquid in the 
beaker, three results can be obtained. 



1. If the liquid in the beaker is alcohol, ether, chloroform, benzene 
or turpentine (to select a few from thousands of similar liquids) the 
ammeter will show no current passing, whatever increase may be 
made in the voltage used, and no chemical action will occur at the 
electrodes. These substances are called non-electrolytes. 

If the liquid is alcohol, 

eiher, chloroform or 


the current will 
NOT pass. 




If the liquid is a solution 

of any acid, alkali or salt 

in water, 

the current will pass and 
decomposition will occur. 

Fig. 38. 
Experiment to show the difference between an electrolyte and a non-electrolyte 

2. If the liquid in the beaker is an aqueous solution of any mineral 
acid, caustic alkali or salt, the ammeter will indicate the passage of a 
large current and chemical decomposition will take place at the elec- 
trodes. These substances are called strong electrolytes. 

3. If the liquid in the beaker is pure water and a very high voltage 
is applied, a sensitive ammeter will register a current but it will be 
very small and the corresponding chemical decomposition at the 
electrodes will be very slight. There also exist a large number of com- 
pounds which, if dissolved in the water, will increase its conductivity, 
to a very small extent. Water, and substances of this type, are called 
weak electrolytes. 




An electrolyte is a compound in solution or a molten compound 
which will conduct electric current with decomposition at the elec- 
trodes as it does so. 

A non-electrolyte is a solution or a molten compound which cannot 
be decomposed by an electric current. 

The electrodes are two poles of carbon or metal by which the 
current enters or leaves an electrolyte. 

The anode is the positive electrode by which the current enters. 

The cathode is the negative electrode by which the current leaves. 


To account for the phenomena of electrolysis the Ionic Theory was 
put forward by Arrhenius about 1880. 

Electrolytes. These substances are believed to contain electrically 
charged particles called ions. 

Ions are derived from atoms (or groups of atoms) but differ from 
them by possessing electrical charges. These charges are positive for 
hydrogen ions and ions derived from metals (or metallic groups like 
NH*), and negative for ions derived from non-metals or acidic radicals. 
The number of electrical charges carried by an ion is equal to the 
valency of the corresponding atom or group. 

Some examples of ionisation are: 


Sulphuric acid 
Sodium chloride 
Sodium hydroxide 
Copper sulphate 
Lead nitrate 
Hydrochloric acid 

It is very important to notice that an ion is very different from the 
corresponding atom, as was explained in Chapter 8. A metallic ion 
is formed from the atom by loss of a number of electrons equal to the 
valency of the metal. Similarly, a non-metallic ion is formed from 
the corresponding atom by gain of a number of electrons equal to the 
valency of the atom, e.g., 

2H + 

so« 2 - 

Na + 


Na + 


Cu 2 + 

so 4 2 - 


2N0 3 - 

H + 


K — e~ — ► K + (univalent) 
Ca - 2e~ — »■ Ca 2+ (divalent) 

iCl, + e-->Ci- 
S + 2e--*-S g - 



These electronic changes give to the ions properties quite different 
from those of the corresponding (electrically neutral) atoms. 

For example, by dissolving ordinary, electrically neutral, molecular 
chlorine in water, a solution is produced which is yellow in colour 
and is a vigorous bleaching agent, but a solution containing clilorine 
ions has neither of these properties. Similarly, ordinary metallic 
sodium, made up of neutral sodium atoms, attacks water liberating 
hydrogen, but sodium ions, Na + , have no such action upon water. 
During the change, 

Na + ±Cl 8 -»-Na + Cl- 
the loss of an electron by Na and the gain of an electron by CI causes 
the ionic particles, Na + and CI" to assume a very stable, inert condi- 
tion; hence their inactivity. 

In strong electrolytes, the ionisation is complete. Thus, there exist 
in a solution of common salt no molecules, NaCI, but only ions, Na + 
and Cl~. All strong electrolytes, i.e., salts, the mineral acids and the 
caustic alkalis, are in this state of complete ionisation in dilute 

In weak electrolytes, ionisation is only slight and most of the 
electrolyte exists in solution in the form of unionised molecules; for 
example, in ordinary bench (2M) acetic acid, out of every 1000 mole- 
cules present, four are ionised and 996 are unionised. 
CH 3 COOH ^ CH 3 COO- + H + 
A solution of ammonia in water is also a weak electrolyte, contain- 
ing a relatively small proportion of ammonium and hydroxyl ions. 
NH«OH ^ NH< + + OH- 
Most of the organic acids are weak electrolytes, e.g., tartaric, citric 
and carbonic. It is not possible to draw an absolutely sharp dividing 
line between strong and weak electrolytes, e.g., trichloracetic acid is 
more highly ionised than acetic acid but is much less highly ionised 
than hydrochloric acid, and so lies between them in strength. For 
your present purpose, the strong electrolytes are the only group of 
considerable importance. 

Non-electrolytes exist only in the form of molecules and are in- 
capable of ionisation, for example, 

Cane sugar 



C,H 6 OH 


Water as an electrolyte 

Water is an electrolyte but is very weak. 
H s O ^ H+ + OH- 



Exact measurement shows that in pure water, for every molecule of 
water ionised, furnishing one hydrogen ion and one hydroxyl ion, 
there are 600 000 000 molecules of water not ionised. The electrical 
conductivity of water, arising from these quantities of ions, is very 
small, but, even so, it must be clearly borne in mind that water is an 
electrolyte and has a small, but measurable, electrical conductivity. 

Further, if by electrical or chemical action hydrogen or hydroxyl 
ions are removed, more water molecules can ionise. So, while at any 
moment H + and OH - concentrations in water are very small, the 
water is potentially capable of yielding more of either ion as circum- 
stances may demand. 1 

Mechanism of electrolysis. Electrolysis of cone hydrochloric acid 

The apparatus of Fig. 39 is suitable. The products are collected 
over calcium chloride solution because both are insoluble in it. 

When the current is switched on, gas will be found to collect in both 
tubes, which were full of saturated calcium chloride solution at the 
beginning. Equal volumes of gas will collect in the two tubes. 

When sufficient gas has accumulated, disconnect the U-tube after 
switching off the current. 

At the cathode, the gas will be colourless and can be tested by 
applying a light to it when it will burn in air. The gas is hydrogen. 

At the anode, the gas will be pale yellowish-green. This gas can be 
tested by damp litmus paper, which will be bleached. It is chlorine. 

Explanation of electrolysis by the Ionic Theory 

So far we have seen that hydrochloric acid contains hydrogen ions 
and chlorine ions. The concentration of hydroxyl ion from water is 
so small that it plays no significant part in this electrolysis. 

When no current is passing, the ions are wandering aimlessly about 
in the solution (Fig. 40). The electrical circuit is closed and, im- 
mediately, the cathode becomes charged negatively, and the anode 
positively. The cathode attracts to itself the positive ions (that is, the 
hydrogen ions) whilst the anode attracts the negative ions (that is, the 
chlorine ions). A procession begins, hydrogen ions to the negative 
pole, chlorine ions to the positive pole. 

1 The exact position is as follows: 

In pure water, 

[H+] = [OH-1 = 10-' mole/dm* (at 25T) 

.•. [H+HOH-] = 10-" = Ktt> (a constant). 
Ktv is called the ionic product of water and is always maintained in aqueous 

If H+ or OH - is withdrawn, water will ionise further to restore the value of 
Kk>. When [H+l = fOH"J = 10"', the liquid is in the same condition as pure 
water, i.e., neutral; if [H + l is greater than 10-', it is acidic; if [H+] is less than 
10-', it is alkaline. But the product [H+l x [OH"] is always 10-". 



l fcarbon) 


solution of 


Fio. 39. 


© i 

© © 


Catnode — 

Ions of hydrogen and 

chlorine wandering 

freely through the 


+ Anode 

•0 «© ©H 

©.^ «! 
-0 ©►©» 


The arrows indicate 
the direction of move- 
ment of the ions when 
charged plates are 
placed in the solution. 


The positive hydrogen ions strike the negative pole and acquire 
from it electrons which make them electrically neutral, and they 
become in this way ordinary atoms of hydrogen which link up into 
molecules and become ordinary gaseous hydrogen. 

Note that it is the electron from the cathode which 
neutralises the hydrogen ion and makes the latter become 
an atom instead of an ion. 

Similarly, the chlorine ion, negatively charged, comes into contact 
with the anode, loses its electron and becomes an atom of ordinary 
chlorine. Pairs of these atoms combine and become molecules of 
ordinary gaseous chlorine, which comes off as a greenish gas. Thus 
the process can be summed up : 

Hydrogen chloride yields ions H + and CI - on solution in water; no 
current is passing. When current passes and electrolysis begins, 

H+ + e--»-H 

CI- - e~ -+ CI 

ion atom 

ion atom 



H + H-»-H a 

CI + CI -*■ c\ t 



At cathode 

At anode 

Hydrogen 1 vol. 

Chlorine 1 vol. 

Fio. 40. 

It should be noted here (see also p. 162) that a cathode, as a source 
of electrons, is equivalent to a reducing agent, and that discharge of 
positively charged ions at a cathode is, chemically, a reduction. The 
hydrogen ion, for example, is reduced (by electron gain) to the 
hydrogen atom: H + + e~— »• H. The electron is supplied by the 
cathode. Similarly, discharge of metallic ions at a cathode is, in all 
cases, a reduction to the ordinary atomic state of the metal, as: 
Na + + <r -> Na or Al 3+ + 3«r ^ Al 

Correspondingly, an anode, as an electron acceptor, is equivalent 
to an oxidising agent. Discharge of negatively charged ions at an 
anode is an oxidation. For example, chloride ion, Cl~, is oxidised to 
the atom by electron loss, the electron being accepted by the anode, 



When two or more ions of similar charge are present under similar 
conditions in a solution, e.g., H + and Na + , or OH~ and S0 4 *-, one is 



preferentially selected for discharge and the selection of the ion dis- 
charged depends on the following factors: 

1. Position of the metal or group in the electrochemical series 



K + 

Ca a+ 

Na + 

Mg* + 

Zn 2+ 

so 4 «- 

Fe s+ 

no 3 - 

To CATHODE «- Pb 8+ 

Cl- -+ To 

H + 






Consider the arrangement given above which is the same as that of 
the electrochemical series. If all other factors (see below) are constant, 
any ion will be discharged from solution in preference to those above 
it, positive ions at the cathode and negative ions at the anode. For 
example, in caustic soda solution, containing positive ions H + (from 
water) and Na + , H + discharges in preference to Na + ; in copper 
sulphate solution, containing OH- (from water) and S0 4 *- as 
negative ions, OH - is discharged in preference to S0 4 *~. 

The more important ions are printed in heavy type in the table. 

2. Concentration 

Increase of concentration of an ion tends to promote its discharge, 
e.g., in concentrated hydrochloric acid, containing OH" (from water) 
and Cl _ as negative ions, the concentration of Cl _ is overwhelmingly 
the greater of the two. In these circumstances, CI - is discharged in 
preference. But, if the acid is very dilute, some discharge of OH - may 
also occur. 

This is the only case you will meet at present in which the order of 
discharge stated by the electrochemical series is reversed by a con- 
centration effect. The same case arises in the electrolysis of sodium 
chloride solution, because the same anions are involved. 

3. Nature of the electrode 

This factor may sometimes influence the choice of ion for dis- 
charge. The most important contrast is electrolysis of a solution of 
sodium chloride, with mercury cathode and with platinum cathode. 



With platinum cathode, H* is discharged in accordance with the 
order of the electrochemical series, Na + being higher in the series. 
The cathode product is hydrogen gas (see p. 148). 

If a mercury cathode is used, there is the possibility of discharging 
Na + to form sodium amalgam with the mercury. This requires less 




Dilute , 

4 — sulphuric 


F10. 41. 

energy than the discharge of H + to hydrogen gas and so occurs in 
preference, and sodium amalgam is the product (see p. 253). 

The products of electrolysis of some important and typical solu- 
tions will now be considered. 

Electrolysis of dilute sulphuric acid (so-called electrolysis of water) 

The apparatus used is shown in Fig. 41. Both electrodes are 
platinum foil. 


The following ions are present: 



From sulphuric acid 
From water 



migrates to the cathode, gains an 
electron and becomes a hydrogen 

H+ + e--»-H 

Hydrogen atoms combine in 
pairs to give molecules. 

H + H->H S 

Migration of S0 4 *~ to the 
anode and discharge of H + are 
equivalent to decrease of con- 
centration of sulphuric acid. 

Hydrogen, 2 vol. 
Acidity decreasing. 

H + and S0 4 2 - 
H + and OH- 


S0 4 2 ~ and OH~ 


both migrate to the anode, 
where OH~, being lower in the 
E.C. series, is discharged in 
preference to S0 4 2 ", in spite of 
the high concentration of the 


By interaction between the OH 
groups, water and oxygen are 

OH + OH 
O + O 

H 2 + 

o 8 


Discharge of OH - disturbs the 
ionic equilibrium of water. More 
water ionises to restore it 

H a O ^ H + + OH- 

Excess H + so produced, with in- 
coming S0 4 8 -, is equivalent to 
increased concentration of sul- 
phuric acid. 

Oxygen, 1 vol. 
Acidity increasing 

The total acidity at anode and cathode together remains constant. 
The final result, 2 vol. of hydrogen at the cathode and 1 vol. of 
oxygen at the anode, is equivalent to the electrolysis of water. 

Electrolysis of caustic soda (sodium hydroxide) solution 

The apparatus is the same as for the electrolysis of dilute sulphuric 
acid (p. 145). The electrodes are again platinum foil. 
The ions present are : 

From sodium hydroxide 
From water 

H + 



Na + 
H + 


both ions migrate to the cathode. 
H + , being lower in the E.C. 
series, is discharged in preference 
to Na\ in spite of the high con- 
centration of the latter. 

H + + e~ ■ 




ions migrate to the anode and 
discharge by loss of an electron. 



By interaction between the OH 
groups, water and oxygen are 

OH + OH 
O + O 

H 2 + O 

The hydrogen atoms then com- 
bine in pairs to give molecules. 

H + H-»-H 2 

Discharge of H + disturbs the 
ionic equilibrium of water. More 
water ionises to restore it. 

H.O ^ H + + OH- 

Excess OH - so produced, with 
incoming Na + , is equivalent to 
increase of concentration of 
sodium hydroxide. 

Hydrogen, 2 vol. 
Alkalinity increasing. 

The total alkalinity at anode and cathode together is constant. 
The process is equivalent to the electrolysis of water. 

Migration of Na + to the cathode 
and discharge of OH - are 
equivalent to fall of concen- 
tration of sodium hydroxide. 

Oxygen, 1 vol. 
Alkalinity decreasing 





Electrolysis of sodium chloride solution 

The apparatus is the same as for concentrated hydrochloric acid 
(p. 142) or dilute sulphuric acid (p. 145). The cathode may be platinum 
(or carbon), but the anode must be carbon to resist attack by chlorine. 

The ions present are: 

From sodium chloride Na + 
From water H + 


Na + 






both migrate to the anode. Cl~ 
is discharged because present in 
much greater concentration than 
OH- (see p. 144). 


The atoms then combine 
pairs to give molecules. 

ci + a -»■ a, 

Chlorine, 1 vol. 


both migrate to cathode. H + , 
being lower in the E.C. series, 
discharges in preference to Na + . 

H + + <r-»-H 

Hydrogen molecules are then 
formed by combination of the 
atoms in pairs. 

H + H -► H, 

Discharge of H + disturbs the 
ionic equilibrium of water. More 
water ionises to restore it 

H,0 ** H+ + OH- 

Excess OH - so produced, with 
incoming Na + , is equivalent to 
the presence of sodium hy- 


Hydrogen, 1 vol. 

Solution becomes alkaline by 

presence of sodium hydroxide. 

Hydrogen and chlorine are produced in equal volumes. If the 
three-limbed voltameter is used, chlorine will have to saturate the 
brine first. 

Electrolysis of copper sulphate solution 

The ions present are: 

From copper sulphate Cu 2+ S0 4 *~ 

From water 



Cu a+ 
H + 

both migrate to the cathode. 
Cu 8+ , being lower in the E.C. 
series, discharges in preference 
to H + . 

Cu* + + 2e" — »- Cu 
The copper deposits as a brown 

Copper deposited. 


I *T Platinum 

Copper copper sulphate 
pfate solution 

Fio. 42. 

Electrolysis of 
copper sulphate solution. 


so 4 2 - 


Platinum anode 

See the exactly similar case of 
dilute sulphuric acid (p. 145). 

Oxygen given off. 
Solution becomes 
sulphuric acid. 

acidic with 

Copper anode 

With this anode, there are three 

1. Discharge of S0 4 8 -] by 

2. Discharge of OH - loss of 

3. Conversion of Cu elec- 
atom to Cu' + . trons 

The last of these occurs most 
readily. S0 4 *- and OH" are 
not discharged. Copper passes 
into solution from the anode as 
Cu 2+ ions. 

Cu-2e-->-Cu ^, - 


Copper passes into solution as 
ions. The total concentration of 
the solution in S0 4 2 ~ (not dis- 
charged) and Cu 2+ (copper is 
depositing on the cathode) is 
constant. The electrolysis merely 
transfers copper from anode to 




Cathode of 

Anode of 

At Cathode 

At Anode 

acid (cone.) 

Carbon or 


one volume 

one volume 

Sulphuric acid 



2 volumes; 
decrease of 

I volume; 
increase of 

Caustic soda 



2 volumes; 
increase of 

1 volume; 
decrease of 

Common salt 

or carbon 


1 volume; 

1 volume 

sulphate (1) 




Oxygen and 
sulphuric acid 

sulphate (2) 





Laws of electrolysis 

The laws expressing the quantitative results of electrolysis were 
first stated by Faraday. They assert that the weight of an element 
liberated during electrolysis is proportional to 

(1) the time of passing the steady current 

(2) the magnitude of the steady current passed 

(3) the chemical equivalent of the element. 

The third statement is the most difficult to understand. It asserts, for 
example, that if, by passing a certain quantity of electricity through 
different electrolytes, hydrogen, copper, silver, chlorine and oxygen are 
given, then, if the weight of hydrogen liberated is l.OO&r g, the weight of 
copper will be 31.8* g, of silver 108x g, of chlorine 35.5* g, and of oxygen 
8a: g, these figures being proportional to the respective chemical equivalents 
of the elements. 

If a current of one ampere flows for one second the quantity of electricity 

conveyed is called one 'coulomb'. It has been found that the passage of 

1 008 
one coulomb of electricity liberates r^= g of hydrogen. So 1.008 g of 

96 500 



hydrogen is liberated by passing 96 500 coulombs ; this amount of electricity 
is called one 'faraday'. It follows from Faraday's Laws that one faraday 
of electricity also liberates the equivalent weight in grams of any other 

Another way of expressing Faraday's Law (3) as stated above avoids 
the use of the term 'equivalent'. The reasoning is as follows. In the case 
of a univalent ion, such as Na+ or CI", one electron is involved in its 
discharge to form an atom. 

Na+ + e~ — *■ Na or CI" — »■ iCU + «" 
From tnis, it follows that one mole of atoms of the element is obtained by 
the passage of one faraday of electricity, i.e., the Avogadro constant, 
6.02 x 10*", of electrons. If the element is divalent (or trivalent), one-half 
(or one-third) of a mole of atoms of it is obtained by expending this 
number of electrons, e.g., 

iZn' 1 " + e~ -> }Zn or iAl»+ +<?-—»• *A1 
Faraday's Law (3) then takes the following form: 

One faraday of electricity liberates, in electrolysis, either one 
mole of atoms of an element or a simple submultiple of one mole. 
One faraday of electricity is often referred to as a mole of electrons because 
it is able to liberate one mole of atoms (or one gram-atom) of a univalent 
metal during electrolysis. (For electroplating, see p. 496.) 

Galvanic (or voltaic) couple 

A galvanic couple is an arrangement of two metals and an electrolyte 
which can generate electric current. To operate effectively the two metals 
must be widely separated in the electrochemical series, e.g., zinc and copper. 
If rods of these metals are placed in the electrolyte, dilute sulphuric acid, 
and the metals are pure, nothing appreciable will occur so long as the metals 
are not in contact. If, however, they are joined by an electrical conductor 
such as a metallic wire (not in contact with the electrolyte), it will be found 
that current will flow in the wire from copper ( + ) to zinc (-), the zinc 
will begin to dissolve in the electrolyte and hydrogen will appear as bubbles 
on the copper. Zinc becomes a cathode and copper an anode in this voltaic 

To produce this result, zinc, the more electropositive of the two metals, 
ionises by electron loss and the electrons pass from zinc to copper through 
the wire. This is equivalent to the flow of conventional current in the 
opposite direction. At the copper surface, the electrons reduce hydrogen 
ions from the electrolyte. 

Zn— »-Zn ,+ + 2e~ 

2H+ + 2e~ — *■ H, 

Theoretically, current can continue to flow as long as materials last; in 

practice, bubbles of hydrogen adhere to the copper, cut off much of its 

contact with the electrolyte and so 'polarise' the cell. 

A cell of this kind is a device for converting chemical energy into 
electrical energy. As the above equations show, the whole process in the 
cell corresponds to the change: 

Zn + 2H+— >-Zn ,+ + H, 
This occurs when zinc dissolves in dilute sulphuric acid and, normally, 
chemical energy is made available from this change as heat ; in this cell. 



much of the chemical energy is converted to electrical energy as electrons 
flow from zinc to copper through the connecting wire. 

taclanchd cell 

This cell is a practicable form of 'primary' cell for producing electric 
current, i.e., it operates by using up the chemicals of which it is composed 
and it cannot be recharged. 

The cathode is a zinc rod or sheet; the anode is a carbon rod and the 
electrolyte in which they are immersed is ammonium chloride solution. 
When the cell is operating, zinc ionises by electron loss and zinc ions 
dissolve in the electrolyte; the electrons pass round the external circuit, 
performing useful work (e.g., lighting a small bulb or ringing a bell), and 
are absorbed at the anode. This causes discharge of ammonium ions from 
the electrolyte, with formation of ammonia (in solution) and hydrogen. 
At cathode At anode 

Zn — >• Zn«+ + 2e~ 2NH«+ + 2e~ — *■ 2NH, + H, 

The anode is immersed in a porous pot containing manganese(lV) oxide 
to oxidise the hydrogen. Otherwise, bubbles of hydrogen adhering to the 
anode would polarise the cell. 

2Mn 4+ + H, + 20H-— »-2Mn s + + 2H,0 

In the so-called 'dry" Leclanche cell, the electrolyte is gelatinised to prevent 
spilling and the cathode is usually sheet zinc, also acting as the cell case. 
This cell will yield a small current continuously without serious polarisa- 
tion. Much larger currents can be obtained intermittently, intervals being 
required to allow the cell to recover from polarisation. (For the lead 
accumulator, see p. 189). 


1. What do you understand by the term 'electrolyte'? Describe experi- 
ments to demonstrate the products formed in the electrolysis of solutions 
of (a) sulphuric acid; (b) sodium sulphate; (c) copper sulphate. (O. and C.) 

2. State Faraday's Laws of Electrolysis. Describe carefully what happens 
when copper sulphate solution is electrolysed between (a) platinum and 
(b) copper electrodes, and when sodium chloride solution is electrolysed 
between (a) platinum, and (b) carbon electrodes. (L.) 

3. Give a general but concise account of the phenomena which occur 
when a salt is dissolved in water and the solution is electrolysed. 

Describe briefly two instances of the practical application of electrolysis. 

4. When an aqueous solution of sodium nitrate is electrolysed with inert 
electrodes, the products are hydrogen (2 vol.) at the cathode and oxygen 
(1 vol.) at the anode. Also, the cathodic liquid becomes alkaline and the 
anodic liquid acidic. Explain these results and write ionic equations in 

5. Explain why solid sodium chloride is a very poor electrical conductor 
while, if melted, it conducts electric current readily. State the products of 
electrolysis of molten sodium chloride and give ionic equations to account 
for them. State and explain the different products obtained by electrolysis 
of a concentrated solution of sodium chloride in water. 

Chapter 13 

Types of Reactions 


Reversible Reactions 
Thermal Dissociation 
Oxidation and Reduction 
Double Decomposition 

CHEMICAL reactions are recognised by certain phenomena 
which accompany them. (See Chapter I.) They can again be 
subdivided into classes of reactions, each of which has its own 


This takes place when two or more substances combine to form a 
single substance. 

Example 1. Iron and sulphur combine when heated to form ferrous 
sulphide (iron(II) sulphide). 

Fe + S -*■ FeS 
Example 2. If warm lead dioxide is lowered into a gas-jar of 
sulphur dioxide, the two compounds combine and lead sulphate is 

PbO^ + SOa-^PbSO* 


This occurs when a compound splits up into simpler substances. 
This change usually lakes place without the necessity for the presence 
of a second substance, and very often the action of heat is sufficient 
to cause the reaction to take place. 

Example 1. If calcium carbonate (for example, marble) is heated 






in an open crucible to bright red heat, the calcium carbonate decom- 
poses into calcium oxide (lime) and carbon dioxide. 
CaC0 3 -*■ CaO + CO, 

Example!. If potassium chlorate is heated strongly it decomposes 
into potassium chloride and oxygen. 

2KC10 3 -»- 2KC1 + 30 a 


This occurs when one element (or group) takes the place of 
another element (or group) in a compound. 

Example 1. If zinc is placed in copper(II) sulphate solution, copper 
is displaced by the zinc and zinc sulphate is left in solution. 

Zn + CuSO t ->- ZnS0 4 + Cu 
Example 2. If chlorine is bubbled into potassium bromide solution, 
the chlorine displaces bromine and a red bubble of bromine is formed. 
A solution of potassium chloride is left 

2KBr + Cl t -»-2KCl-fBr 1 
Reversible reactions 

Suppose that n molecules of a substance A can react with m mole- 
cules of a substance B to produce x and y molecules respectively of 
products C and D, the system being homogeneous (i.e., entirely 
liquid or entirely gaseous). 

nA + mB ^xC + yD 
As soon as a little of C and D is formed, a reverse reaction will begin. 
At first the forward reaction will preponderate but, as C and D 
accumulate, the reverse reaction will build up until an equilibrium 
position is reached with forward and reverse reactions proceeding 
at the same rate. The composition of the mixture will then appear 
constant, though it is the net result of the two opposing reactions, 
not a static situation. This kind of chemical behaviour (which is 
very common) is called a reversible reaction. Several examples of 
reversible reactions are considered a little later in this chapter. 

Rate of reaction. The rate of a chemical reaction (that is, the 
amount of reaction occurring in unit time) is influenced by several 
factors such as temperature, pressure (if gas is involved), concentration 
of reagents, light and catalysis. 

Temperature. In general, rise of temperature tends to increase the 
rate of a chemical reaction by two factors: molecules of the reactants 
move more rapidly at higher temperature, producing more frequent 
and more energetic collisions. For example, a mixture of hydrogen 
and oxygen (2:1 by volume) remains unchanged indefinitely at 

ordinary temperature but the introduction of a local source of high 
temperature, e.g., an electric spark, starts an explosive combination 
of the two gases. 

Concentration. If the reagents involved in a chemical reaction have 
their concentrations increased in a homogeneous mixture, more 
frequent molecular collisions will occur and the reaction rate will 
be thereby increased. A simple case of this occurs when addition of 
some concentrated acid increases the rate of liberation of hydrogen 
by the action of zinc on dilute hydrochloric acid. In the case of mixed 
gases, increase of pressure implies increase of concentration and 
tends to increase the rate of reaction between the gases because of 
more frequent molecular collisions. Reactions in the liquid or solid 
state are very little influenced by pressure because liquids and solids 
are almost incompressible. 

If the reacting system is not homogeneous, e.g., if a solid is reacting 
with a liquid, particle size may influence reaction rate. For example, 
aluminium foil reacts moderately with bench sodium hydroxide 
solution when warmed but powdered aluminium reacts rapidly from 
cold and will usually froth out of the test-tube spontaneously. 

2A1 + 2NaOH + 6H 2 -► 2NaAl(OH)< + 3H 2 
This occurs because, for a given weight of metal, powder offers a 
much greater area to the reacting liquid than foil. In the limit, sub- 
division to actual atoms or molecules (such as occurs in solution) 
offers maximum opportunity for reaction. 

Light is a source of energy and can influence the rate of some chem- 
ical reactions considerably by energising some of the molecules 
involved. For example, the reaction between chlorine and hydrogen 
at ordinary pressure is negligible in darkness, slow in daylight but 
explosive in sunlight (at room temperature). Also, light is vital to 
the photosynthetic production of starch by plants from carbon di- 
oxide and water (p. 186). 


Catalysis is said to occur when the rate of a chemical reaction is 
altered by an agent (the catalyst) which is left unchanged in amount 
and in chemical nature at the end of the reaction. 

A catalyst usually increases the rate of a reaction and this is 
called positive catalysis. It is found that, in a reversible reaction, a 
given catalyst influences the rate of the forward and reverse reactions 
equally. It has no influence on the actual equilibrium position which 
is reached; it only enables the equilibrium to be attained more 
rapidly. The equilibrium position is determined by factors such as 
concentration of reactants, temperature and pressure, not by the 
catalyst used. 



In general, a catalyst will function even though it is present in only 
minute proportion, e.g., the rate of decomposition of hydrogen per- 
oxide is measurably increased by the introduction of one ten-millionth 
of its weight of finely divided platinum. 

A catalyst may change its physical nature during a reaction, e.g., 
coarsely powdered manganese dioxide may become fine powder (in 
the first example below) but it must be left chemically unchanged at 
the end of the reaction. 

Examples of positive catalysis 

1. Heating of potassium chlorate 

2KC10 3 -> 2KC1 + 30 2 

2. Synthesis of sulphur trioxide 

2SO a + O, ^ 2SO, 

3. Synthesis of ammonia 

N 2 + 3H 2 ^ 2NH 8 

4. Decomposition of hydrogen 

2H 2 2 -»-2H 2 + 0, 


Manganese dioxide, Mn0 2 
(Manganese(IV) oxide) 

Vanadium pentoxide, V.0 6 
or platinum (powder) 

Reduced iron (powder) 

Manganese dioxide, Mn0 2 
or platinum powder 

A negative catalyst decreases the rate of a chemical reaction. This 
variety of catalysis is occasionally used to suppress an unwanted 
reaction (e.g., 2% of alcohol in chloroform acts as negative catalyst 
for the oxidation of the chloroform to poisonous products by the 
air) but has obviously fewer uses than positive catalysis. 

Chemical equilibrium 

It has already been mentioned (p. 154) that, in suitable conditions, 
a reversible reaction will reach an equilibrium in which forward and 
reverse reactions are proceeding at equal rates and the composition 
of the mixture is kept constant. The conditions affecting chemical 
equilibrium will now be considered in relation to a few reversible 
reactions of industrial or experimental importance. In the consider- 
ation of such reactions, a very valuable contribution can be made by 
a rule known as Le Chatelier's Principle. It is very widely applied to 
physical and chemical situations and, for the latter, takes the fol- 
lowing form. 

If a chemical system is in equilibrium and one of the 
factors involved in the equilibrium is altered, the equili- 
brium will shift so as to tend to annul the effect of the 



Consider the important industrial conversion of sulphur dioxide 
to sulphur trioxide, a stage in the Contact Process for the manufac- 
ture of sulphuric acid. 

2S0 2 + 0,^ 2S0 3 (all reagents gaseous) 

From left to right, this reaction is exothermic, i.e., it liberates heat. 
Suppose that an equilibrium position has been reached at certain 
temperature, pressure and concentrations of reactants. 

Effect of temperature change. Suppose the temperature of the 
reacting system to be lowered (with no other change). Le Chatelier's 
Principle requires the system to react so as to oppose this change, 
i.e., to raise the temperature again. To do this, heat must be liber- 
ated by causing more sulphur trioxide to be produced, i.e., lower 
temperature is favourable to production of sulphur trioxide. At once 
the difficulty arises that lowering of temperature reduces the rate 
of reaction so increasing the time required for the production of 
sulphur trioxide. For this reason a catalyst must be introduced, 
e.g., vanadium pentoxide, V 2 5 , which increases the rate of reaction 
and makes it viable at a relatively low temperature (450°C) which is 
favourable to the production of sulphur trioxide. In general, low 
temperature gives an equilibrium favourable to an exothermic 
reaction but catalysis is needed to give a favourable reaction 

It is obvious that, conversely, high temperature produces an 
equilibrium favourable to an endothermic reaction. In this case, the 
rate of reaction is also improved by the rise of temperature so a 
catalyst is not needed in principle. For example, in a now obsolete 
industrial process for making nitric acid, nitrogen and oxygen were 
combined to form nitrogen monoxide (an endothermic reaction) 
simply by passing them through an extended electric arc (temper- 
ature around 2200°C). 

N»(g) + 0,(g) ^ 2NO(g) (L -»- R 180.6 kJ absorbed) 

Concentration of reactants. Suppose that, in the Contact Process 
reaction above, equilibrium has been reached in certain conditions 
and then the concentration of oxygen is raised (relative to sulphur 
dioxide). Le Chatelier's Principle requires the system to react to 
oppose this change, i.e., to reduce the oxygen concentration towards 
its former level. This can only be done by combining it with sulphur 
dioxide to form sulphur trioxide, i.e., increased concentration of 
oxygen favours conversion of more sulphur dioxide to trioxide. 
(Correspondingly, increased concentration of sulphur dioxide 
favours conversion of more oxygen to sulphur trioxide.) In practice, 
the sulphur dioxide and air mixture used contains about three times 
as much oxygen as is theoretically required for the sulphur dioxide 



content. (Use of more air renders the product too dilute in sulphur 

Pressure. Pressure changes are significant only in reactions con- 
cerning gases because liquids and solids are almost incompressible. 
Taking the Contact Process reaction again, suppose that, in a system 
in equilibrium in the gaseous state, pressure is increased (with no 
other change). Le Chatelier's Principle requires the system to react 
to oppose the change, i.e., to reduce the pressure towards its former 
value. Since 2 molecules of sulphur dioxide and one molecule of 
oxygen produce 2 molecules of sulphur trioxide, a total 3 volumes of 
sulphur dioxide and oxygen are converted to 2 volumes of sulphur 
trioxide (Avogadro's Law, T. and P. constant). That is, the system 
must convert more sulphur dioxide and oxygen to sulphur trioxide 
to reduce pressure by reducing volume. In general, if a given (entirely 
gaseous) reaction proceeds with a reduction in the number of mole- 
cules present, it is favoured by high pressure. (Conversely, gaseous 
reactions involving increase in the number of molecules present are 
favoured by low pressure.) In practice, the Contact Process can reach 
a satisfactory 98% yield of sulphur trioxide (calculated on the S0 2 
used) without recourse to pressure above atmospheric but the similar 
reaction for producing ammonia, 

N 2 + 3H 2 v± 2NH 3 (4 vol -»■ 2 vol, T. and P. constant) 

uses 400 atm pressure or more. See also p. 431. 

Reaction of iron with steam 

This reaction is usually performed by boiling water in a flask and 
passing steam over iron filings in an iron tube at bright red heat 
(p. 265). It produces hydrogen and tri-iron tetroxide and the reaction 
is reversible. 

3Fe(c) + 4H 2 0(g) «* Fe 3 4 (c) + 4H 2 (g) 

In the conditions stated above, the reaction goes almost to com- 
pletion from left to right in the sense that the iron can be converted 
almost entirely to the oxide. This is so because incoming steam main- 
tains its concentration at a high level (forcing the reaction from left 
to right) while hydrogen is swept out of the reaction tube and its 
concentration continually tends towards zero (keeping the reverse 
action slight). 

If a current of hydrogen is passed over red-hot trifcrric tetroxide, 
the conditions are reversed and the oxide can be almost completely 
reduced to iron. 

If the iron tube is sealed, an equilibrium will ultimately be set up 
with all four materials present and the forward and reverse reactions 
maintaining the equilibrium. The composition of the mixture at 



equilibrium is determined mainly by the temperature employed. 
Pressure is without influence on the equilibrium position because 
there is no change in the number of gaseous molecules present as a 
result of the reaction. 

Reaction of water with bismuth chloride 

This reaction gives visible evidence of reversibility. If dilute hydro- 
chloric acid is added carefully, with shaking, to a gram or two of 
bismuth carbonate, a colourless solution will ultimately be obtained 
containing bismuth chloride. If this solution is put into a large 
beaker and water is carefully added with stirring, the liquid soon 
becomes turbid with the formation of a fine white precipitate of 
bismuth oxychloride. 

BiCl 3 + H t O ^ BiOCl + 2HC1 
This occurs because addition of water increases its relative concen- 
tration and forces the reversible reaction from left to right (as written 
above). In terms of Le Chatelier's Principle the system reacts so as 
to tend to reduce the increased relative concentration of water. If 
concentrated hydrochloric acid is then added drop by drop with 
stirring, the turbidity can be cleared as the increased concentration 
of hydrochloric acid forces the reaction from right to left (or, in terms 
of Le Chatelier's Principle, as the system reacts to reduce the in- 
creased concentration of hydrochloric acid). By alternate careful 
additions of water and hydrochloric acid, the reaction can be shown 
to reverse itself several times (by alternate turbidity and clearance) 
until the dilution becomes too extreme for a precipitate to be seen. 

A thermal dissociation is a reversible reaction brought about by the 
application of heat. Examples are the thermal dissociations of 
ammonium chloride and dinitrogen tetroxide. 

NH 4 Cl^NH a + HCl; N 2 4 ^2N0 2 
In both cases, complete dissociation doubles the number of mole- 
cules and the volume (T. and P. constant). Consequently, since the 
mass of material is unchanged and the volume is doubled, the vapour 
density of the product is ultimately halved. For example, at about 
350°C, the vapour density of ammonium chloride is 14.5 against a 
required value of iNH 4 Cl or 26.75 for NH 4 C1 molecules. This 
represents about 85% dissociation. Similarly, the vapour density of 
dinitrogen tetroxide which is close to 46 just above its boiling-point 
(22°C), corresponding to molecules N 2 4 (92), gradually falls with 
temperature rise until, at about 150°C, it is about 23, corresponding 
to N0 2 (46) molecules. At intermediate temperatures, mixtures of 
N a 4 and NO a molecules are in equilibrium and give vapour densities 
between 46 and 23. The gas also changes colour from pale yellow 



through reddish brown to almost black as temperature rises, and 
reverses the colours with cooling. 

The above data are for one atmosphere pressure. If an equili- 
brium position exists at a certain temperature and pressure and then 
the pressure is increased, Le Chatelier's principle requires both 
systems to oppose this change, i.e., reduce pressure towards its 
former value. This they can do by reducing the degree of dissociation. 
This reduces the number of molecules present and so reduces volume 
and pressure. Thus, in both cases, dissociation is reduced by in- 
creased pressure at constant temperature; conversely, reduced 
pressure increases dissociation. 

Thermal decomposition is the name given to the break-up of a 
compound by heat without any recombination on cooling. An 
example is the conversion of ammonium nitrite to water and nitro- 
gen by heat. 

NH 4 N0 8 -»- 2H a O + N, 

Double decomposition 

This name has been given to reactions in which two compounds 
take part, both are decomposed and two new substances formed by 
an exchange of radicals. Double decomposition reactions are always 
of the type: 

A.B + C.D-»-A.D + C.B 
For example: Cu.S0 4 + H 8 .S -> Cu.S + H 8 .S0 4 

Commonly, both die original compounds used in the reaction are 
soluble in water, while, of the products formed, one (sulphuric acid) 
is soluble and one (copper sulphide) is not. Usually the precipitated 
compound is the one which is wanted, for it can easily be separated 
and purified by filtration and washing. Less frequently, the im- 
portant product of a double decomposition reaction is more volatile 
than the other compounds concerned and is driven off either as a gas 
or, by heating, as the vapour of a volatile liquid. For example: 
NaCl + H 8 SO« -*- NaHS0 4 + HC1 f 

KN0 3 + H 8 S0 4 -► KHSQ 4 



It must be observed, however, that, from the modern point of 
view, many reactions of 'double decomposition', especially those 
occurring in solution, are regarded as taking place between com- 
pounds which are already fully ionised, so that no decomposition 
takes place. The situation is merely that, if the ions present in the 
mixture can form an insoluble combination, they will do so, and the 



corresponding compound will precipitate. For example, if a solution 
containing sodium chloride (that is, the ions Na + and CI") is mixed 
with one containing silver nitrate (that is, the ions Ag + and NO a _ ), 
silver chloride precipitates. The essential change is represented in the 
form: Ag + CI - — *■ AgCl 4. . Ions Na + and NO„- remain in solu- 
tion. Such reactions of ion aggregation are virtually irreversible 
because the very low solubility of the precipitate suppresses any 
possible reverse with the dissolved ions. 

Oxidation and reduction 

It is necessary to discuss the way in which the use of these terms has 
developed before it is possible to give a concise definition of them. 

Originally, reduction was a reaction in which some compound was 
deprived of all, or part, of the oxygen it contained, and an oxidation 
was a reaction in which a substance combined with oxygen. In the 
reaction represented by the equation: 

CuO-f- H 8 -»-Cu + H 8 

the hydrogen is oxidised to water and the copper oxide reduced to 
copper. This is the simplest possible use of the terms. 

The readiness with which hydrogen combines with oxygen to form 
water caused hydrogen to be regarded as a kind of 'chemical opposite' 
of oxygen; so the term oxidation was extended to include reactions in 
which a compound gave up some or all of its hydrogen as well as 
those in which it combined with oxygen. (This idea of oxidation is 
analogous to the idea of enriching a man by relieving him of his 
debts.) By this extension a reaction like the conversion of con- 
centrated hydrochloric acid to chlorine is called an oxidation because 
two molecules of hydrogen chloride (2HC1) are converted to a mole- 
cule of chlorine (Cl 8 ) by loss of hydrogen. 

2HCl + (0)->H 8 + Cl» 
The oxygen for this reaction is commonly supplied from manganese 
dioxide (p. 355). A similar case is the oxidation of hydrogen sulphide 
to sulphur by, say, chlorine. 

H 2 S + Cl 8 -*2HCl-r-S 

A further extension of the idea of oxidation arose from the fact that 
certain elements exercise two or more different valencies. One of the 
most important of these is iron, which exerts a valency of 2 in iron(II) 
compounds and 3 in iron(III) compounds. Now the conversion of 
iron(II) oxide, FeO, to iron(UI) oxide, Fe 8 3 , is a clear and simple 
case of oxidation. 

2FeO + (O) -*■ Fe 8 O a 



But all iron(ll) salts, for example, iron(II) chloride, FeCl 8 , and 
iron(II) sulphate, FeS0 4 , correspond to, and can theoretically be 
obtained by neutralisation from, iron(II) oxide; similarly, all iron(III) 
salts, for example, iron(III) chloride, FeCl 3 , or iron(III) sulphate, 
Fe 2 (S0 4 ) 3 , correspond to, and can theoretically be obtained by 
neutralisation from, ferric oxide. So the term oxidation was ex- 
tended to include not only the conversion of iron(II) oxide to 
iron(III) oxide, but also the conversion of any iron(Il) compound to 
an iron(III) compound. Thus, the conversion of iron(II) chloride to 
iron(lll) chloride by chlorine 

2FeCl, + Cl a -»-2FeCla 

is spoken of as being an oxidation, although oxygen is not involved. 
Similar cases of oxidation are: 

Fe t (S0 4 ) 3 + H 2 

2Fe(OH) 3 


2FeSO« + H„S0 4 + O ■ 

2Fe(OH) 2 + H 1 0- r -0- 

In such cases it will be noted that the oxidation always involves an 
increase in the proportion of those groups which are electronegative, 
that is, which migrate to the anode when the compound is electrolysed 
(p. 144). Thus, there is a higher proportion of the electronegative 
Cl~ group in iron(III) chloride, FeCl 3 , than in iron(II) chloride, FeCl s , 
and a higher proportion of the electronegative S(V - group in iron(III) 
sulphate, Fe 2 (S04) 3 , than in iron(H) sulphate, FeS0 4 . To summarise 
the above development, we can say, so far, that oxidation has in- 
cluded the following ideas : 

1. Combination with oxygen. 

2. Removal of hydrogen. 

3. Increase in the proportion of electronegative ion in a metallic 
compound (which requires increase in the valency of the metal). 

Reduction is the reverse effect in each case. 

Recently these ideas have been restated in electronic terms as 
below, with considerable gain in simplicity in most cases. From the 
electronic point of view, the following definitions can be given: 

Oxidation is the process of electron loss. 
Reduction is the process of electron gain. 
An oxidising agent is an acceptor of electrons. 
A reducing agent is a donor of electrons. 

Oxidation and reduction always occur together; they are com- 
plementary processes of electron loss and electron gain respectively 



and must occur simultaneously. The electrons lost by the reducing 
agent must be accepted by the oxidising agent present. The following 
examples illustrate these ideas. 

1. When magnesium is oxidised by combination with oxygen, the 
metal is oxidised by losing two electrons per atom. These electrons 
are accepted by oxygen atoms, which are reduced as a result. Mag- 
nesium (giving out electrons) is the reducing agent; oxygen (accept- 
ing electrons) is the oxidising agent. 

Mg-2e--»-Mg s+ ; *O t + 2e~ -»- O*- 

Magnesium oxide is a collection of Mg 2+ and O a ~ ions in equal 

Notice that the combination of magnesium with chlorine or sulphur 
is a similar process, as: 

Mg + Cl 2 -*■ Mg* + .2C1- 
Mg + S -* Mg* + .S*- 

Chlorine and sulphur must, like oxygen, rank as oxidising agents; 
as before, magnesium is a reducing agent in these reactions. From 
this point of view, any conversion of a metal to its ions is oxidation, 
i.e., electron loss; correspondingly, any conversion of a non-metal to 
its ions is reduction, i.e., electron gain. 

2. If a metallic ion, e.g., the iron(Il) ion, Fe*+, is so treated that 
it loses a further electron, it is oxidised and is a reducing agent. The 
process is: 

Fe* + - e~ -> Fe 3 + 

An iron(lll) ion is formed. An agent must be present, e.g., chlorine, to 
accept the electrons made available by the ferrous ions. It acts as the 
oxidising agent (electron acceptor) and is reduced. 

iCl, + e — ► Cl- 

The complete reaction can be represented: 

Fe* + + iCl t -»• Fe 3+ + Cl~ 

It will be observed that the valency of the metal increases from 2 to 3 
during the oxidation. 

3. The 'removal of hydrogen' aspect of oxidation is interpreted in 
the following way. Consider the oxidation of hydrogen sulphide by 
chlorine. Hydrogen sulphide is slightly ionised as: 

H a S ^ 2H+ + S 2 - 

The sulphide ion parts with its two electrons and is, therefore, 
oxidised, acting as a reducing agent. 

S a ~ - 2e- -*- S 


The electrons are accepted by chlorine atoms, so that chlorine acts as 
an oxidising agent and is reduced. 

Cl 2 + 2e~ -»- 2C1- 

Adding the two equations, S 2_ + Cl 2 — *■ S + 2C1 - . The hydrogen 
ion of the hydrogen sulphide is unchanged. 

4. The reduction of hot copper(Il) oxide by hydrogen is rendered 

Cu a+ .O s - + H 2 -*- Cu + H a O 

It is clear that the copper(II) ion is reduced by electron gain, as: 

Cu*+ + 2e~ -*■ Cu 

The two electrons are made available by the reaction between the 
oxide ion, O 2- , and hydrogen: 

O a - + H 2 -*• H 2 + 2e~ 
By combining with oxygen in this way and supplying electrons to the 
metallic ion, hydrogen exercises reducing properties. The oxide ion 
is oxidised by electron loss and the oxygen atom remains in com- 
bination with hydrogen as water. 

5. Nitric acid can operate as an oxidising agent, by accepting 
electrons, in several different ways. The two of greatest importance 

4HN0 3 + 2e~ ->- 2N0 3 " + 2H t O + 2N0 2 (i) 

8HN0 3 + 6e~ -> 6N0 3 " + 4H a O + 2NO (ii) 

The electrons are supplied by a reducing agent. A metal often acts in 
this way, e.g., 

Cu ->• Cu*+ + 2e~ 

To supply the two electrons needed in equation (i), one copper atom 
is required. This yields the reaction : 

Cu -f 4HN0 3 -> Cu 2+ .(N0 3 -) 2 + 2H 8 + 2N0 2 

This is the chief reaction occurring when copper reacts with concen- 
trated nitric acid. The products are copperfjl) nitrate, water and 
nitrogen dioxide. 

To supply the six electrons needed in reaction (ii), three copper 
atoms are required. This yields the reaction: 

3Cu + 8HNO a — >• 3(Cu 2+ .2N0 8 ~) + 4H 3 + 2NO 

This is the principal reaction when copper reacts with a mixture of 
concentrated nitric acid and water in equal volumes, and is the 
recognised laboratory preparation of nitrogen monoxide, NO. This 
gas is also the product when iron(II) sulphate solution is warmed 
with nitric acid of suitable concentration : 

6Fe a+ + 8HNO a -»■ 6Fe 3 * + 6NO a " + 4H 2 Q + 2NO 



The iron(II) ions are oxidised (by electron loss) to the iron(IIl) state. 
Below are mentioned some common oxidising and reducing agents 
and their usual mode of operation in terms of electron exchange: 

Oxidising agents 
Oxygen (p. 277) 
i 2 + 2e-->0*~ 

Reducing agents 
Hydrogen sulphide (p. 398) 
H 2 S ^ 2H + + S*- 

S 2 - -» S + 2e~ 

Chlorine (p. 362) 
iCl 2 + e--*-Cl- 

Ozonc (p. 283) 

3 + 2e-->0 2 + 2 - 

Sulphur dioxide (aqueous) (p. 404) 
S0 2 + H 2 ^ H 2 SO a ^ 2H + + S0 3 *- 
S0 3 *- + H s O ->• S(V" + 2H + + 2e~ 

Hydrogen (with heated metallic oxides) 

(p. 267) 
0*--r-H 2 ->-H 2 + 2e- 

Hydrogen peroxide (p. 280) Carbon monoxide (with heated metal- 

H 2 8 + 2H + + 2e— >2H 2 lie oxides) (p. 301) 

The H" is supplied by water 0*~ + CO — »- CO a + 2e~ 

or dilute acid present in the 


Nitric acid (p. 438) Carbon (with heated metallic oxides) 

(p. 290) 
4HNO, + 2e--> 0»" + C -> CO + 2e~ 

2NO," + 2H 2 + 2N0 2 or 20»- + C — CO s + 4e" 

8HNO, + 6e~ -*■ t , . ,._. 

6NO - + 4H 2 + 2NO (For oxidation number, see p. 5Z/J 

Notice that all the oxidising agents are electron-acceptors; all the 
reducing agents are electron-donors. Examples of their oxidising or 
reducing action will be found on the pages quoted. 

Exothermic and endothermic reactions 

The great majority of chemical reactions are accompanied by a 
marked heat change. The above two types are distinguished. 

An exothermic reaction is one during which heat is liberated to the 

e.g., the burning of hydrogen or carbon 
H t (g) + 40,(g) -* H t O(l); A/7 = -286 kJ (-68.4 kcal) g-eqn" 1 

C(c) + ojffi -* CO^g); A// = -406 kJ (-97 kcal) g-eqn"* 

Heat energy can be derived from coal, coke, coal-gas, petrol and 
paraffin, in the home or in industry, because the combustions of all 
these substances are exothermic reactions. 



An cndothermic reaction is one during which heat is absorbed from 
the surroundings. • 

e.g., formation of nitrogen monoxide or carbon disulphide from 
their elements. 

*N 2 (g) + iOafe)-*" NO(g); AH = +90.3 kJ (+21.6 kcal) g-ean"' 

C(c) + 2S(c) ->- CS 2 (1) ; AH = + 1 17 kJ (+28 kcal) g-eqn" > 
The symbol, AH, denotes the heat change taking place in the reaction 
corresponding to the equation to which it is attached (with gram units 
of weight, i.e., g-equation -1 , and with stated physical condition of 
reagents at constant pressure). The usual convention now used is to 
refer the heat change to the reacting system, i.e., if heat is lost by 
the reacting system (an exothermic reaction), AH is negative; if heat 
is taken in by the reacting system (an endothermic reaction), AH is 

A kcal is a kilocalorie and is the heat required to raise the tempera- 
ture of one kilogram of water by 1°C. This is one thousand times as 
large as the ordinary calorie which raises the temperature of one 
gram of water by 1°C. J is the symbol for the joule, which is now 
replacing the calorie as the international unit of energy. The joule is 
expressed in electrical terms; thus, if one coulomb of electricity passes 
at an electrical pressure of one volt, the energy involved is one joule. 
In general, 

volts x coulombs = joules 
The relation between the calorie and the joule is: 

1 calorie = 4. 18 joule 
The symbol, kJ, represents 1000 joule. 

The following sections state in outline how some common therrao- 
chemical estimations can be made. 

Heat of combustion 

For this determination, a bomb calorimeter is suitable. It is usually 
made of steel (for strength), nickel-plated on the outside and pro- 
tected from oxidation by a coating of enamel or (better) gold leaf on 
the inside. A known weight (a g) of a compound is placed in the 
platinum cup, C. Air is displaced by oxygen which is then allowed to 
reach 20-25 atm pressure and the bomb is closed by a screw valve. 

The bomb is immersed in a known quantity of water in a well- 
lagged calorimeter fitted with stirrer and accurate thermometer. Elec- 
tric current (through insulated platinum leads) heats the iron wire, T, 
which fires the experimental compound. This causes a sudden rise of 
temperature which is read on the thermometer (or, better, expressed 
on a temperature-time curve with readings at 15 sec intervals). If the 


water-equivalent of the whole system (bomb 
calorimeter, water, thermometer, calori- 
meter, stirrer) is W and the temperature 
rise is t°C, the total evolution of heat is Wt 
cal. If the heated iron wire yields x cal in 
burning, the heat of combustion of a g of 
material is (Wt — x) cal. The heat of com- 
bustion of one mole of compound is then 

{Wt — x) x — cal, where M is the molecular 

weight of the compound. (This estimation is 
at constant volume since any change in the 
capacity of the bomb calorimeter is negli- 
gible.) To convert to joules, multiply the 
above quantity by 4.18. 


Fig. 43. 
Bomb calorimeter. 

Heat of neutralisation, solution and precipitation 

Heat of neutralisation can be determined by the apparatus shown 
in Fig. 44. The apparatus is allowed to stand to attain a steady 
temperature (which is read on the accurate thermometer). The test- 
tube is then broken, the acid alkali solutions mix and are stirred and 
there is a temperature rise which is read on the thermometer. (Better, 
the temperature can be read at 15 sec intervals and the results 
rendered as a graph.) There are present 100 cm 3 of solutions and 
dilute solutions can be considered as having the same heat capacity 
as water. If the water-equivalent of the whole apparatus is W g, and 
the temperature rise is t"C, the heat evolved by the reaction is 
(100 + W)t cal. Since one-twentieth of a mole of acid and alkali is 
involved, the heat of neutralisation is (100 + W)t X 20 cal. (Multi- 
ply by 4. 1 8 for joules.) As explained on p. 226, the heat of neutralisa- 
tion of any strong acid by any strong alkali is a constant (13.7 kcal 
or 57.3 kJ). 

Heat of solution can be determined in the apparatus of Fig. 44, with 
a known weight, a g, of finely powdered solute in the inner tube and 
a known quantity of water (b cm 3 ) in the vacuum flask. The steady 
temperature is read and the inner tube is broken. The solid is dis- 
solved with stirring and the final (usually lower) temperature is read. 
The heat change is then (b + W)t cal, where Wis the water equivalent 
of the apparatus and / is the observed temperature change. This is 
the heat of dilution of the solute to a concentration of a g in b g or 
cm 3 of water. It is also the heat of solution if the dilution is so great 
that further dilution produces no further heat change, and if it is 
referred to one mole of solute. 



Heat of precipitation can be determined in the same way as heat of 
neutralisation, using any pair of solutions that can precipitate (com- 
pletely for practical purposes) an 'insoluble' product, e.g., M solu- 
tions of silver nitrate and sodium chloride. If 50 cm 3 of each are used, 
the experiment and calculation is the same as for the neutralisation. 
Ag'(aq) + Cl-(aq) -»- AgCI(c); A// = -65.7 kJ g-eqn"> 




Vacuum flask 

50cm 3 of M NaOH 
50cm 3 of M HCI 

Fig. 44. 
Heat of neutralisation. 

Conservation of energy. All energy changes which occur during 
chemical and physical changes of the kind considered above must 
conform to the Law of Conservation of Energy; that is, energy can 
only be changed from one form into its equivalent of another form 
with no total loss or gain. 

To take the combustion of sucrose (cane sugar) to carbon dioxide 
and water as products (at ordinary temperature and pressure), we 
have the equation 

CmHmO^c) + 120 2 (g)-> 12CO s (g) + 11H.OG); 

AH = -5685 kJ or -1360 kcal g-eqn" 1 

During this reaction, covalent bonds between carbon, hydrogen and 

oxygen atoms in the sugar molecules are replaced by covalent bonds 



between carbon and oxygen atoms in carbon dioxide and between 
hydrogen and oxygen atoms in water. These changes involve altera- 
tions in electron orbits and, therefore, energy changes. Also, the 
forces holding sugar molecules together in crystals are overcome and 
replaced by forces holding water molecules together as liquid. At the 
same time, molecules of carbon dioxide are left as gas and molecules 
of water as liquid, both types of molecule possessing energy of 
motion. If these varying energy changes (and any others we may have 
forgotten) arc allotted their correct signs and quantities, and account 
is similarly taken of the heat energy liberated during the combustion, 
the total energy change (obtained by adding these quantities to- 
gether) must be zero. 

The dissolution of a solid in water may be a purely physical change, 
i.e., the separation of ions or molecules from crystals and their dis- 
persal into the water. This requires a supply of energy to overcome 
inter-ionic forces (as in a Na+Cl" lattice) or van der Waals forces (as 
m many organic solids). This energy is taken from the water (as heat) 
and the temperature falls, i.e., dissolution of a solid in water is often 
endothermic. In some cases, however, chemical reaction occurs, e.g., 
ions may be hydrated, as with anhydrous copper(II) sulphate 

Cu 2+ + 4H.O -* (Cu.4H a O) 2+ 
In such a case, this heat of hydration (added to the true heat of solu- 
tion) may make the total change exothermic. 


1. Give an account of catalysis, briefly referring to the catalytic oxida- 
tion of ammonia. Describe an experiment to show that the manganese 
dioxide used in the ordinary laboratory preparation of oxygen from 
potassium chlorate has not changed appreciably in weight during the 

2. What is a reducing agent ? Give three examples of common reducing 
agents. Describe and explain any experiment in which sulphuric acid is 
reduced. (O. and C.) 

3. Describe and explain experiments in which the following substances 
play the part of oxidising agents: (a) nitric acid; (b) copper oxide; (c) 
chlorine. How would you show practically that oxidation has occurred in 
two of these cases you select? (O. and C.) 

4. Define reversible reaction, thermal decomposition and thermal dis- 
sociation. Describe any experiment you have seen to demonstrate thermal 

What happens when (a) lead nitrate; (6) mercuric oxide; (c) nitrogen 
dioxide; (d) ammonium nitrate are heated? In each case state to which of 
the above classes the reaction belongs, giving your reasons. (N.U.J.B.) 

5. State what type of reaction — combination, double decomposition, 
replacement, oxidation — takes place when : (a) sulphur is boiled with nitric 
acid; (fc) ammonia gas and hydrogen chloride are mixed; (c) chlorine is 
passed into potassium iodide solution; (a 1 ) hydrogen sulphide is passed into 



a solution of copper sulphate; (e) excess of carbon dioxide is passed into 
lime-water. (N.U.LB.) 

6. Explain and illustrate by one example in each case the meaning of 
the terms: (a) reversible reaction; (6) catalyst; (c) allotropy; (d) double 
decomposition; (e) deliquescence; (/) synthesis. (L.) 

'•, , N 1 (g) + 3H 1 (g)^2NH > (g) 

l ne large-scale manufacture of ammonia by the above reaction is catalysed 
by iron. Select from the following possibilities the correct explanation of 
its catalytic activity; the iron: (1) increases the rate of the reaction left to 
right only, (2) decreases the rate of the reaction right to left only, (3) in- 
creases the rate of both forward and reverse reactions, (4) produces an 
equilibrium mixture which is more concentrated with respect to ammonia. 

8. CH 3 COOH(l) + C,H,OH(l) ^ CH,COOC,H,(l) + H.Ofl) 

This system (entirely liquid) will come to equilibrium at room temperature 
and pressure in several hours, starting from one mole each of acetic acid 
and ethanol. There is then two-thirds of a mole of ethyl acetate (and of 
water) present. The reaction is catalysed by hydroxonium ion. State two 
practical methods by which you would expect to obtain the equilibrium 
mixture almost unchanged in composition but much more rapidly. State, 
and briefly explain, one method by which you would secure the conversion 
of more than two-thirds of the acetic acid to ethyl acetate. State and 
briefly explain the (qualitative) effect of adding more water to the equili- 
brium mixture. What, if any, would be the effect on the equilibrium mixture 
of raising the pressure to 2 atmospheres? 

9. The reaction: 3Fe + 4H.O ^FejO, + 4H„ is reversible in con- 
ditions in which iron and its oxide are solid and the other reagents gaseous. 
Describe and explain what occurs (a) if hydrogen is passed over red-hot 
tri-iron tetroxide in an open tube, (i) if steam is passed over red-hot iron 
in an open tube. If an equilibrium system exists among all these four 
reagents in a closed vessel at red heat, what will be the (qualitative) effect, if 
any, on the equilibrium of (c) adding more of the oxide of iron, (d) doubling 
the total pressure, (e) doubling the partial pressure of hydrogen, tempera- 
ture being constant throughout ? Briefly explain your reasoning in each case. 

10. The reaction: N, + O s ^ 2NO, is reversible and (from left to right) 
endothermic, all reagents being gaseous. Explain the term endothermic 
reaction. If the above system is in equilibrium (at a temperature which 
allows quite rapid reaction), what, if any, will be the effect on the equili- 
brium of (a) doubling the total pressure, (6) doubling the partial pressure 
of nitrogen, (c) lowering the temperature slightly? Explain briefly why this 
reaction (when used industrially) was not catalysed. What alternative was 

1 1 . For each of the following, state one change of conditions (tempera- 
ture changes excluded) suitable for increasing the rate of reaction : (a) com- 
bination of hydrogen and chlorine, (/>) decomposition of hydrogen peroxide 
solution, (c) displacement of hydrogen from hydrochloric acid by mag- 
nesium. Hydrochloric acid is a stronger acid than sulphuric acid. What 
explanation in terms of chemical equilibrium can you give of the following 
facts: (a) addition of concentrated sulphuric acid to sodium chloride 
liberates hydrogen chloride quantitatively, (b) addition of dilute sulphuric 
acid precipitates barium sulphate quantitatively from barium chloride 

Chapter 14 

Air, Combustion and Rusting 

A STUDY of the air begins naturally with an examination of that 
most familiar of all chemical reactions— burning. Most of the 
common combustible materials (coal, wood, petrol) are complex com- 
pounds which are unsuitable as the starting-point of our investigation. 
We shall fall back upon the metals, which are all elements, and, there- 
fore, the simplest materials known. 

Effect of heating certain metals in air 

Copper. Take up a piece of copper foil in a pair of tongs and hold 
it in a Bunsen flame. The metal becomes red hot and, on cooling, is 
covered with a black layer. This is black copper oxide. If the metal is 
scraped, the surface layer is obtained as a powder and the fresh copper 
exposed can be similarly treated. In time, a quantity of the black 
powder can be obtained. . 

Lead. Heat a little lead foil on a crucible lid. It melts to shining 
beads of molten lead. Stir the beads. The metal gradually changes to 
a yellow powder called litharge. 

Magnesium. Hold one end of a length of magnesium ribbon in 
tongs and place the other end in a Bunsen flame. The ribbon burns 
with a dazzling flame (it is rather dangerous to look at it for any 
length of time) and leaves a white ash— magnesium oxide. 

These experiments leave no doubt that the metals concerned have 
undergone a drastic change. The products left after heating them are 
quite different from the original metals. 

The nature of the change 

It is hardly necessary to describe experiments to prove that air is 
concerned in this change. This can, however, quite readily be shown 
by taking the most combustible of these metals, magnesium, placing 
it in a crucible, filling the crucible with sand, well pressed down to 
exclude air, and heating the crucible to redness. On cooling, the 
magnesium is unchanged. 





Since air is concerned in the change, two possibilities have to be 
considered. The metals may have combined with something from 
the air or they may have lost something, which has been taken up by 
the air. (A third possibility is that the change in the metals may be 
due to some rearrangement of their material without loss or gain, but 
this is unlikely since such a change could presumably occur without 
air.) Clearly, in the first case, the material which had combined would 
make the weight of the product greater than that of the metal, while, 
in the second case, the material lost would have the reverse effect. We 
have only to weigh the metal before burning and the product after 
burning to decide this point. 

To find whether there is any change in weight when magnesium burns in 

Weigh a crucible (with lid) con- 
taining about 0.5 g of magnesium. 
Set up the apparatus as in Fig. 45. 
Remove the lid and heat the crucible. 
When the magnesium begins to burn, 
put the lid on the crucible. Raise it 
occasionally to allow air to enter to 
burn the magnesium but, as far as 
possible, avoid losing any 'smoke' 
(fine particles of magnesium oxide) 
which would tend to make the final 
weight too low. When all the 
magnesium has burned allow the 
crucible and lid to cool. Then weigh 
them again. There will be a gain in weight. 

The products of combustion of a substance always weigh more than 
the original substance. 

This is even true for coal. If we could collect all the ash, soot 
smoke and the gases (carbon dioxide and steam) which escape up the 
chimney, they would weigh more than the original coal. (See p. 25.) 
This gain in weight, which occurs no matter what the substance is 
that burns, at once establishes the point that, during burning, the 
burning material combines with something. We now have to show 
whether the 'something' comes from the air. 

To do this, it will be necessary to find whether, when a substance 
burns in air, the air decreases in amount. We must devise an experi- 
ment to test this point. It would be absurd to carry out the experiment 
in the open laboratory into which air could leak from outside to 
replace loss; we must secure a sample of air in a confined space, that 

Fig. 45. 

Burning magnesium in a 



j s in a closed vessel, and, if possible, any change in the amount of 
air should be automatically shown to us by the apparatus. A very 
simple arrangement satisfies this requirement. We shall confine the 
sample of air in a bell-jar over water (see Fig. 46). The water forms 
a flexible base to the bell-jar and will move up or down inside the jar 
to show us what is happening to the amount of air inside. As a matter 
of mere convenience we shall choose yellow phosphorus for our 
burning substance this time. (Yellow phosphorus takes fire very 
readily and must be treated with great care. Never touch it with your 
fingers. The heat of them may start it burning and the burns it will 
cause are very severe and difficult to heal. Yellow phosphorus is 
always kept under water because of the ease with which it takes lire.) 

To find whether there is a diminution in the volume of the air when 
phosphorus burns in it 

Float a small porcelain dish on water in a pneumatic trough and 
put in it a piece of yellow phosphorus about as big as a pea. Place 
over it a bell-jar and adjust the water to level A (Fig. 46, I), the 
stopper of the bell-jar being removed. The bell-jar above A is 

As A 

fhosphorai „ » 

Phosphorus just begin- After the phosphorus 
ning to burn has burnt 

Fio. 46. 

The burning of phosphorus in air. 

graduated into five equal portions. Heat a long iron needle in a 
Bunscn flame, touch the phosphorus with it, withdraw the needle 
quickly and insert the stopper of the bell-jar. The phosphorus burns 
with a bright yellow flame, giving off dense white fumes of phos- 
phorus pentoxide, which fill the jar. After a time the phosphorus no 
longer burns. The water level inside the bell-jar will then be found to 



rise 1 and, when the bell-jar is cold, the water level will stand at mark 
B on the jar (Fig. 46, II). Clearly, the rise of water inside the bell-jar 
means that, during the burning of the phosphorus, some of the air 
was used up to combine with the phosphorus. 

After a time, the white fumes dissolve in the water, leaving the bell- 
jar clear. It will then be seen that some unburnt phosphorus remains 
in the porcelain dish. This is very significant. The Same was not 
extinguished for lack of phosphorus. We can see that some gas still 
remains in the bell-jar from the mark B upwards. This gas must be 
different from ordinary air because it will not allow phosphorus to 
burn in it; it must also be different from the part of the air which has 
combined with the phosphorus because it will not do this. If a 
lighted splint Is plunged into the residual gas the splint is extinguished. 
We are forced to conclude, therefore, that the air is not a single sub- 
stance. It must contain at least two gases — one which supports the 
combustion of phosphorus and one which does not. Further, we may 
conclude that the gas which is active in supporting the combustion of 
the phosphorus constitutes about one-fifth of the air by volume (this 
represents the rise of the water from A to B) and the other gas about 
four-fifths. These two gases have names. The one which supports 
the combustion of phosphorus is called oxygen, the other nitrogen. 

Let us sum up in a few sentences what we have learnt so far. 

The principal gases in air are oxygen and nitrogen. 
Oxygen constitutes about one-fifth of the air by volume and 
nitrogen about four-fifths. During the combustion of a sub- 
stance, it combines chemically with the oxygen of the air 
and the chemical combination is accompanied by the 
evolution of light and heat. The combination with oxygen 
causes a gain in weight. Nitrogen will not support com- 

By an experiment similar to the above, it can be shown that the 
material of a burning candle combines with oxygen and causes the 
water-level inside the bell-jar to rise. The candle will not, however, 
remove all the oxygen. 

We may now give the equations for the chemical reactions con- 
sidered in this chapter. 

r ■! W . h j lethe Phosphorus is burning, the level of water inside the bell-jar will 
fall. This is an expansion effect of the heated air. It is also necessary to pour 
water into the trough until the levels of water inside and outside the bell-jar are 
equal. If this is omitted we are not measuring the volumes under the same 
conditions. ^^ 


2Cu + 0,-)-2CuO 

copper(Ii) oxide 

2Pb + O a ->2PbO 
lead(II) oxide 

2Mg+ 2 ->-2MgO 

P« +5O 8 -*P 4 O 10 

The smouldering of phosphorus 

Phosphorus smoulders in air. The chemical effects of the smoulder- 
ing are very similar to those of the active burning of phosphorus 

f \ Yellow 


\no longer 

Phosphorus starting After several days 
to smoulder 

Fio. 47. 
Smouldering of phosphorus in an enclosed space. 

except for the time factor. This can be shown by an experiment for 
which Fig. 47 is sufficient explanation. 

More accurate determination of the proportion of oxygen in air by 

In this experiment, we make use of the smouldering of phosphorus 
to absorb the oxygen from a measured volume of air. (See also above.) 

Take a graduated glass tube, closed at one end, fill it to a depth 
of about 5 cm with water, close the open end with the thumb and 
invert the tube in a deep jar of water. (If possible, allow the tube 
to stand like this for several hours so that the air is saturated with 
water-vapour.) Adjust the level of water in the graduated tube to be 











beginning to 



After the phos- 
phorus has ceased 
to smoulder 

Fio. 48. 

To find the percentage by 
volume of oxygen in the air. 

the same as the level in the jar. 
The air inside the tube is then at 
atmospheric pressure. Read off 
the volume of air. Now push up 
inside the tube a flexible wire carry- 
ing a piece of yellow phosphorus 
(Fig. 48, i). Read the temperature 
of the laboratory and the barometer 
and set the apparatus aside until the 
phosphorus no longer smoulders. 
The water level inside the tube will 
have risen to C, the remaining gas 
being nitrogen (Fig. 48, ii). Remove 
the phosphorus and then lower 
the graduated tube until C is at 
the level of water in the jar, giving 
atmospheric pressure again inside 
the tube. Read off the volume of 
nitrogen. Take the temperature 
of the laboratory and read the 
The calculation of the percent- 

age of oxygen in the air by volume is given below. 

= 70.5 cm 5 
Pressure 755 mm 

= 55.0 cm 3 
Pressure 760 mm 

Original volume of air 

Temperature 14°C 
Final volume of nitrogen 

Temperature 12°C 

The volume of nitrogen must first be converted to the volume it 
would occupy at the same temperature and pressure as that of the 
original air. 

Volume of nitrogen at I4°C and 755 mm pressure 

= 55.8 cm 3 
/. volume of oxygen = (70.5 - 55.8) cm 3 

= 14.7 cm 3 
.'. percentage of oxygen in the air by volume 
_ 14.7 
= 20.8 

X 100 

In dry air, the correct percentage of oxygen is 20.9 by volume. 


to Suction 

Other gases present in air 

Carbon dioxide. Carbon dioxide is present in air to the extent of 
0.03% by volume. It is formed during the combustion of all the 
common fuels— coal, coke, coal-gas, water-gas, petrol, paraffin oil- 
all of which contain carbon. 

c + o»-*co 2 

It is also breathed out as a waste product by all animals. 

In spite of the enormous quantities poured into the atmosphere in 
this way, the proportion of it remains constant, partly because carbon 
dioxide is taken up by the leaves of plants and converted to complex 
starchy compounds (for a more complete discussion of this subject 
see p. 185) and partly because it dissolves in the water of the oceans. 

The presence of carbon dioxide in the 
air can be shown by aspirating air 
through a boiling-tube containing a 
little lime-water. After a time the lime- 
water will go turbid, showing the pre- 
sence of carbon dioxide (Fig. 49). 

Ca(OH) 2 4- CO a -> CaCO s + H 2 

Water-vapour. This substance is 
always present in the air in varying 
quantities. It is given off by evapora- 
tion from the oceans, rivers and lakes. 

Its presence may be demonstrated by 
exposing some deliquescent substance, 
say calcium chloride, to the air on a 
clock-glass. A solution of the com- 
pound will be obtained after a day or 

two. If this is distilled (see p. 191), the colourless liquid obtained 
may be proved to be water by the tests given on pp. 270-1. 

The rare gases. About 1 % of the air by volume is made up of the 
rare gases. The most abundant of them is argon, the others being 
neon, xenon, krypton and helium. The proportion of each of the last 
four is very minute. Argon and neon have found a use in 'gas-filled' 
electric light bulbs and coloured 'neon' electrical signs. They are 
obtained from liquid air. 

Impurities. The air always contains small traces of many gases- 
hydrogen sulphide, sulphur dioxide and others— especially in in- 
dustrial areas. They are given off during the combustion of coal and 
fuels derived from it. The tarnishing of silver is chiefly due to the 
formation of a layer of black silver sulphide on it by the action of 
traces of atmospheric hydrogen sulphide. 

\s water 

Fio. 49. 

Method of showing the 

presence of carbon dioxide 

in the atmosphere. 


The following diagram summarises the composition of the air. 














by volume 

by volume 



about 1% 


by volume 

by volume 


Exposure of some compounds to ordinary air 

Sodium hydroxide. If exposed to air, sodium hydroxide absorbs 
water-vapour from the air and forms a solution; that is, deliquescence 
occurs (p. 524). No chemical action is involved so far. The solution 
then absorbs carbon dioxide from the air and forms a crystalline 
solid, washing soda or sodium carbonate decahydrate. 

2NaOH + CO, + 9H s O ->- Na s CO 3 .I0H 2 O 
If left to stand, this solid will eventually lose nine of its ten molecules 
of water of crystallisation (per molecule) to the air. That is, the 
decahydrate effloresces, falling to a fine white powder consisting of 
small crystals of the monohydrate, Na 8 C0 3 .H,0. Some sodium 
bicarbonate may also be formed as the monohydrate absorbs carbon 
dioxide from the air. 

Na a C0 3 .H a O + CO a -> 2NaHC0 3 

Calcium chloride deliquesces in air and forms a solution. This 
tendency to absorb water-vapour explains the use of calcium chloride 
as a drying agent for gases (not ammonia, with which it combines). 
In this context, however, the relative amount of moisture is small 
and the anhydrous calcium chloride is only hydrated to the solid 
hydrate, CaCI 3 .6H a O. This hydrate is also deliquescent. 

Phosphorus pentoxide deliquesces in air, forming a colourless solu- 
tion containing metaphosphoric acid. 

P 4 O 10 + 2H t O-*-4HPO 3 

The pentoxide is also used as a drying agent for gases, but not for 
alkaline gases like ammonia. 

Concentrated sulphuric acid absorbs water from the air, diluting 
itself, usually up to about three times the original volume. The acid 
is said to be hygroscopic (not deliquescent, this term being reserved 
for solids which absorb water from the air to form solutions). 



The efflorescence of washing soda was noticed above. Another 
efflorescent compound is Glauber's salt, sodium sulphate decahydrate, 
Na,SO 4 .10H a O. On exposure to air it loses the whole of its water of 

Iron(II) sulphate heptahydrate, FeS0 4 .7H a O, is also efflorescent, i.e., 
loses water of crystallisation on exposure to air. In addition, it under- 
goes oxidation by oxygen of the air, acquiring brown patches of basic 
iron(M) sulphate. 

12FeS0 4 + 6H a O + 30, ->• 4{ Fe 8 (S0 4 ) 3 .Fe(OH) 3 } 

Composition of air by weight 

This estimation was carried out by Dumas, 1841. Air was drawn 
through the following apparatus in the order shown (Fig. 50). 

Heated copp er 

Fig. 50. 
Gravimetric composition of the air. 

1. U-tubes containing potassium hydroxide to remove carbon 
dioxide (only one shown in the figure). 

2. U-tubes containing concentrated sulphuric acid (on pumice) 
remove water-vapour (only one shown in the figure). 

3. A heated, weighed tube containing finely divided copper to 
absorb oxygen; 

and, finally, the remaining 'atmospheric nitrogen* (still containing 
the rare gases) entered a weighed evacuated globe. 

The increase in weight of the copper gave the weight of oxygen, 
and the increase in weight of the globe, the weight of nitrogen (and 
rare gases). 

Neglecting carbon dioxide, the percentage of oxygen by weight in 
dry, pure air is 23.2% the remainder being nitrogen and rare gases. 



Air— mixture or compound? 

The evidence on which a decision on this question can be reached 
is given below: 

(1) The composition of air is very nearly, but not quite, constant. 
Small, but definite, differences of composition have been de- 
tected when samples of air from different parts of the earth 
have been analysed. 

(2) (a) If air is dissolved in water and boiled out again (see p. 193), 
the percentage by volume of oxygen in the air boiled out is 
increased from 21% to about 30%. No chemical reaction is 
involved here; the composition of air has been altered by a 
physical method, which depends merely on the fact that oxy- 
gen is twice as soluble in water as is nitrogen. 

(b) If liquid air is allowed to evaporate, nitrogen evaporates 
more quickly, leaving almost pure oxygen. Here again, the 
gases of air are separated by a purely physical process. 

This evidence alone is sufficient to decide that air is a mixture. 
Confirming it are the following facts: 

(3) If nitrogen, oxygen, carbon dioxide, water-vapour and the 
rare gases are mixed in appropriate proportions, there is no 
explosion, evolution of heat, volume change or other evidence 
of chemical combination, but the product resembles ordinary 
air in every way. 

(4) The composition of air corresponds to no simple chemical 
formula such as it would be expected to possess if it were a 

The combustion of a candle 

The products of combustion of a candle can be shown by the 
apparatus of Fig. 51. 

The products from the burning candle are drawn up the funnel and 
through the apparatus. 

Drops of liquid will condense in the U-tube and the anhydrous 
copper sulphate will change to blue hydrated crystals. This proves 
that one of the products when a candle is burned is water. The lime- 
water will rapidly turn milky (the mil kiness is caused by a fine preci- 
pitate of chalk). This proves that carbon dioxide is also given off 
from the candle. 

Candle-wax contains carbon and hydrogen. During the burning, 
these elements combine with oxygen of the air, forming carbon 
dioxide and water. 




c + o t - 

2H. + 0,- 
CuS0 4 + 5H,0 

Ca(OH) t + C(V 

► CO, 

► 2H.O 

»- CuS0 4 .5H 2 

hydrated crystals; 

► CaC0 3 + H,0 


to Filter -pump 
or Aspirator 

Lime water 

rstneiti White anhydrous 

Canalt Copper Sdlphate 

Fio. 51. 
To show the products of combustion of a candle. 

Most of the common fuels— coal, coke, coal-gas, petrol, paraffin 
oil, water-gas— contain one or both of the same elements as candle- 
wax and their products of combustion consist mainly of carbon 
dioxide and water. 

Though they pass off as gases, the products of combustion of a 
candle should weigh more than the candle-wax, which has burnt. 

To show that the products of combustion of a candle weigh more than 
the candle-wax burnt 

The apparatus is described by Fig. 52. 

Suspend the apparatus from the balance hook of a large, rough 
balance, and add weights to counterpoise it. Light the candle. The 
apparatus will quickly gain in weight and will depress the pan of the 
balance to which it is attached. 

The gain in weight is the weight of oxygen from the air with which 
the carbon and hydrogen of the candlewax have combined during 

Other kinds of combustion 

Combustions in oxygen or air are so common that it becomes 
almost habitual to use the word 'combustion' as if it referred to this 



kind of reaction alone. Actually, it may be applied to any chemical 
combination accompanied by light and heat in which one or more 
of the reactants are gaseous. You will find, for example, in this book, 
accounts of the combustion of hydrogen, phosphorus and copper in 
chlorine gas. 

Cottori wool 

Ca/aum Chloride 
—to absorb 
water vapour 

h— Cotton wool 

Soda -lime 
to absorb 
Carbon Dioxide 

Wire gauze 


Fig. 52. 
To show the increase in weight when a candle burns. 

Oxygen of the air burning in coal-gas. An interesting reversal of 
the usual state of affairs, in which air is acting as the combustible 
material, may be secured by using apparatus as in Fig. 97. (See 
p. 303.) 


The important facts connected with the rusting of iron may be 
ascertained by the following experiments. 

To find if there is any change in weight when iron rusts 

Weigh a clock-glass containing some iron borings. Damp the bor- 
ings and set them aside to rust. When rusted, place them in an oven 
to dry thoroughly, then weigh the clock-glass again. There will be 
a gain in weight. In this respect rusting is analogous to burning. We 
may now try an experiment to see whether the air is similarly 
concerned in both. 



To find whether iron combines with anything from the air while 


This experiment is described by Fig. 53. 

To show the character of the gas left in the bell-jar in Fig. 53 (II), 
fill up the trough till level B rises to level A. Then remove the 
stopper of the bell-jar and insert a lighted taper. It will be extin- 
guished; the remaining gas is nitrogen. During rusting the iron has 
combined with the oxygen of the air. This accounts for the rise in 
the water level (Fig. 53) and the gain in weight (see p. 182). 


Iron nails 
in a wet 
muslin bag 

vJH&fcrT-i- li 


r/se oF 





Iron about to rust 


After rusting for three or four days 
Fig. 53. 

The rusting of iron. 

Rusting and burning are similar 

It is important to understand clearly that, from the chemical 
standpoint, rusting and burning are the same process. During burn- 
ing, magnesium combines with oxygen of the air, forming magnesium 

2Mg-f-O t -»-2MgO 

During rusting, iron combines with oxygen of the air in the presence 
of water to form brown hydrated ferric oxide, 'rust'. 
4Fe + 30 2 -*2Fe 2 3 

The only difference is in the time required for the two processes. Heat 
is generated during rusting just as it is during burning, but it is dis- 
sipated to the surroundings without attracting notice because of its 
much slower rate of production. 



It is very unfortunate for mankind that iron, which possesses so 
many useful properties, should be so readily attacked by the oxygen 
of the air. To protect iron from rusting, it is painted or galvanised 
(see p. 483). The various protective processes probably cost at least 
fifty million pounds annually over the whole world. 

We will now investigate further the conditions under which iron rusts. 
We have already seen that during rusting, iron combines with oxygen 
of the air, and it is common knowledge that iron rusts more readily 
when plenty of water is present. It will be of interest to separate 
these agents and find their individual effects. To do this, we must 
expose iron to the action of air in the complete absence of water (to 
water-free air) and to water in the complete absence of air (to air- 
free water). 

To find the effect of exposing iron to air and water separately 

(1) Exposure of iron to air separately. To do this, set up apparatus 
as in Fig. 54, 1. The figure sufficiently explains the experiment. 

.Rubber tubing 


to dry the 

Cotton wool 
Iron nails 



Air -tree 

Iron nails 

Fio. 54. 

To show thai iron will not rust in the presence of: 
1, air alone; II, water alone. 

(2) Exposure of iron to water separately. Boil about 350 cm 3 of 
water rapidly in a beaker for at least half an hour. This will boil all 
the air out of it. Put a few iron nails into a test-tube and fill it to the 
brim with the boiled water. Press into the mouth of the test-tube a 
rubber stopper carrying a glass tube and rubber tube as shown in 
Fig. 54, II. (Note that the glass tube must be flush with the bottom 



of the stopper.) The water will rise into the rubber tube. Place a clip 
in position to exclude all air. 

Leave the test-tubes for several days. In neither case will 
rusting occur. This means that iron will not rust in the 
presence of air alone or of water alone; both are needed 
together to rust the iron. 
To make sure that the nails you have used are actually capable of 
rusting, put them into a test-tube with a little water, leave the test- 
tube open to the air and notice the result after a day or two. 

What we have done in the above experiment is to take two sub- 
stances, air and water, which normally act on iron together and test 
the effect of each singly. This is the favourite device of scientists. It 
is only by finding out the effects produced by one agent at a time 
that reliable information can be obtained. 

Oxygen and carbon dioxide in life-processes. Photosynthesis 

It can readily be shown that carbon dioxide is present in the air expelled 
from the lungs of a human being. Use the apparatus of Fig. 55. 

With the clip A open and the clip B closed, breathe in air through the 
mouthpiece M. The air bubbles through the lime-water in C. Then close 
clip A, open clip B and breathe out the air so that it passes through the 
lime-water in D. Repeat this several times. The. lime-water in C remains 
unaffected while that in D is rapidly turned milky. This must mean that 
during its occupation of the lungs, the air has increased its proportion of 
carbon dioxide, which reacts with the lime-water in D to form a precipitate 
of chalk. 

Ca(OH), + C0 3 -*-CaCO s + H.O 

At the same time, oxygen is absorbed from the air into the blood forming 
a loose compound with the haemoglobin (the red colouring matter) of the 
red blood corpuscles. 

The use of the oxygen and the presence of carbon dioxide may be ex- 
plained briefly as follows. The process of digestion converts our food 
materials into compounds which are either soluble in water or easily 
emulsified with it. The soluble compounds are absorbed into the blood as 
the food passes through the small intestine. They are carried round by the 

Fio. 55. 
To show the presence of carbon dioxide in respired air. 



blood-stream and used either to replace wastage and maintain growth in 
the body-tissues or to supply energy for movement of the body and to 
maintain its temperature at about 98.4°F, a temperature usually consider- 
ably higher than that of surrounding objects. This energy is supplied by 
oxidation of the soluble products of digestion. Consider, for example, a 
sugar. It is well known that if a sugar is thrown on to the fire it burns 
vigorously, giving off heat, and forming, as products of combustion, carbon 
dioxide and water. The same oxidation process, giving the same products, 
occurs in the body, the oxygen being taken from its loose combination 
with hemoglobin to oxidise the dissolved sugar. For a given quantity of 
sugar oxidised, the same amount of heat is given out whether it is burnt 
rapidly on the fire or oxidised more slowly in the body. 

QHj.O, + 60, — *■ 6CO, + 6H.O; a definite quantity of heat liberated 

This is the source of the body's heat. The waste product, carbon dioxide, 
is carried round in the blood (chiefly as bicarbonate) and is liberated as the 
blood passes through the lungs, from which it is breathed into the air. The 
breathing process of animals is similar. 

It has already been pointed out that all common fuels give off carbon 
dioxide when burnt. Living animals also discharge carbon dioxide as we 
have just seen. There must obviously be some agency at work using up 
these vast quantities of carbon dioxide and restoring the oxygen absorbed 
from the air during combustion, for otherwise the composition of the air 
would change appreciably. 

Actually, the plant life of the world is at work restoring the balance. With 
the help of energy from the sun and with chlorophyll (the green colouring 
matter of leaves) as catalyst, plants are continually building up carbo- 
hydrates (e.g., starch and sugar) from water and carbon dioxide of the air. 

Carbon dioxide + water 4- energy —*■ starch + oxygen 

Oxygen is liberated and passes into the air. Notice that this process 
(photosynthesis) is endothermic. The energy absorbed is released again when 
carbohydrates are used as food and oxidised in the blood-stream to carbon 
dioxide and water (see p. 187). It is obvious that the waste product, carbon 
dioxide, breathed out by animals and given off by carbonaceous fuels, is 
the raw material from which plants build up carbohydrates. The waste 
product, oxygen, from the plant is the gas essential for animal breathing 
and combustion of fuels. These opposing processes, at work together, keep 
the composition of the air constant. 

In addition to carrying on the above feeding process, plants also breathe 
in the same way as animals, but, relatively, their breathing process is of 
small account. On balance, a plant uses up carbon dioxide and liberates 
oxygen. It may be mentioned that, biologically, the characteristic difference 
between animals and plants is not connected with size or movement or 
similar factors, but lies in the fact that plants build up (synthesise) complex 
starchy food materials from the simple compounds, carbon dioxide and 
water, while animals must have these complex materials available as food 
and break them down by digestion into simpler substances. 

Experiments illustrating the above brief outline will be found described 
in any elementary text-book of biology. 

Extraction of starch 

Starch is a carbohydrate, having a chemical formula of the type C x (H,0) v 
— in this case, C,H 10 O 6 , as the empirical formula. Production of starch by 


photosynthesis (above) and storage of the product is a plant's method of 
providing a reserve of nutrition and energy, e.g., in seeds and tubers. To 
extract starch, potatoes (for example) are pulped with water and the starch 
is washed out. Sieving retains cellulose and allows starch to pass; water is 
then centrifuged away and the starch dried. 

Starch has many uses. Wheat starch is the main constituent of bread in 
Europe and America, and maize starch of cornflour. Potato starch is widely 
consumed as food and other edible forms are sago, tapioca, rice and arrow- 
root. Potato starch is used in the manufacture of glucose, dextrins and 
starch syrup. Until recently starch was the main source of ethanol (alcohol) 
by fermentation and a great quantity of alcoholic beverages is still pro- 
duced this way, but industrial alcohol is now largely made from petroleum 
products such as ethylene (p. 334). 

When used as human food, starch is hydrolysed by enzyme action to 
glucose which is utilised in three principal ways: for production of energy 
by oxidation to carbon dioxide, for storage as fat and for storage as 
glycogen, a compound which readily breaks down to glucose as the body 
needs it. 

Starch can be chemically hydrolysed in several successive stages. Cold, 
dilute mineral acid converts it to soluble starch, which dissolves in hot 
water and does not set to a paste on cooling. Dry heat (as in baking bread) 
produces dextrins of the kind which colour and flavour bread crust or 
toast and more complete hydrolysis produces glucose (a sugar). 

Test for starch. Starch itself, and soluble starch, give a dark blue colora- 
tion with iodine, which disappears on heating but reappears on cooling. 
(The more complex dextrins produce a red colour with iodine but simpler 
dextrins and glucose give no colour reaction with iodine.) 

Conversion of starch to glucose 

Starch is mixed with about 2\ times its weight of water and about 1% 
of sulphuric acid is added. This mixture is heated under pressure for about 
90 min, after which glucose solution is left. 

(C,H 10 O 6 ). + «H a O -► wC,H„0, 

A slight excess of chalk is added to destroy the remaining acid and calcium 
sulphate is filtered off. After decolorisation with charcoal and evaporation 
under reduced pressure, glucose crystallises out after cooling and can be 

Structure of the starch molecule 

A starch molecule is produced by the combination of many molecules 
of glucose, C ( HitO ( , with elimination of molecules of water. Since this 
splitting out of H,0 molecules is classed as condensation and many mole- 
cules of glucose are involved, starch can be called a condensation polymer of 
glucose. For this purpose, glucose can be written as HO.C,H 10 O 4 .OH. 
Hydroxyl groups from a pair of different glucose molecules eliminate a 
water molecule between them and this happens repeatedly to form a very 
large linear molecule in the following way. Letting C ( Hi,0, = X, 

HO-X-o lHTHOr -X— Q (H + HOr -X-O lH + HO| — X— OH -> 

HO— X— O— X— O— X— O— X— OH + 3H 2 



Between 200 and 500 glucose molecules can group themselves together 
in an actual starch molecule, giving a possible molecular weight in the 
region of 60000. If the chain-molecule formed remains straight, the 
material is called amylose, but if a branched structure is formed with many 
offshoots (usually about 20 glucose units long), the product is amylopectin. 
Various samples of starch have varying proportions of the two. Hydrolysis 
with hot, dilute, sulphuric acid puts back the H,0 molecules and breaks 
up the chain to single glucose molecules again. The two components of 
starch differ somewhat in behaviour with iodine. Amylose gives a blue 
colour and amylopectin reddish purple. 


1. What happens when the following substances are exposed to the air: 
(a) quicklime, (b) solid sodium hydroxide, (c) washing soda, and (rf) 
anhydrous calcium chloride? How would you confirm experimentally your 
answer in one case? (N.U.J.B.) 

2. What do you understand by the term combustion ? Give two examples 
of combustion in which oxygen plays no part. Describe experiments which 
illustrate the resemblance between the combustion of a candle and the 
respiration process in an animal, giving sketches of any apparatus you 
would use. (D.) 

3. An ordinary paraffin candle consists of carbon and hydrogen in com- 
bination. State what becomes of the candle when it burns in air, and 
describe an experiment in support of your statement. Make a sketch of the 
apparatus you would use to show that the weight of the products of com- 
bustion is greater than the weight of the candle consumed. (N.U.J.B.) 

4. Explain concisely the chemical meaning of the terms combustion and 
rusting. Describe the experiments you would make to show what reactions 
take place when iron rusts in moist air. (O. and C.) 

5. When copper is heated in air, a black substance is formed. How 
would you prove that the copper had combined with another element to 
form this new substance, and how could you determine what this element 
is? (N.U.J.B.) 

6. If a green plant is placed in a closed volume of air in daylight, what 
changes would you expect to result in the composition of the air? How 
would you demonstrate that t hese changes occur ? To what do you attribute 
them? (O. and C.) 

7. (a) Give briefly three reasons for regarding air as a mixture and not a 

(6) How would you determine the weight of oxygen in a given 

volume of air? Sketch the apparatus you would use and show, 

using an imaginary case, how you would work out the calculation, 
(c) Describe briefly how you would determine the volume of oxygen 

in a sample of air confined over water in an eudiometer tube. 


8. Describe in detail how you would proceed in order to remove com- 
pletely the oxygen, carbon dioxide, and water-vapour from ordinary air. 
How would you prove that the residual gas was really free from these 
substances? (L.) 



9 How would you find by experiment the percentage, by volume, of 
oxv'een in the air? What other gases are always present in the air? 

What is a physical change? Describe a simple laboratory experiment in 
which some 'air' is obtained of different composition from the ordinary air 
as the result of a physical change. (L.) 

10 State briefly the chemical nature of starch and its function in plants. 
Given a sample of starch, how would you prepare in the laboratory a 
reasonably pure solution of glucose from it? Give a chemical test by which 
vou would demonstrate that this solution no longer contains starch. 
Glucose contains carbon, 40.0%, hydrogen, 6.7%, and oxygen, 53.3%. 
Show that its empirical formula is CH,0. If its molecular weight is 180, 
what is its molecular formula? (C = 12, O = 16, H = 1.) 

Lead accumulator (reference from p. 152) 

The lead accumulator is a secondary cell, i.e., it has to be 'charged' 
by passage of direct current (usually rectified mains current) through it, 
after which it will 'discharge' yielding direct current for use where required, 
e.g., in the electrical system of an automobile. These processes of charge 
and discharge can be repeated many times. 

The cathode and anode of the accumulator are both grids ot leaa- 
antimony alloy. At discharge, both grids carry a lead sulphate filling. The 
electrolyte is sulphuric acid suitably diluted with water. 
During charge, the following changes occur. 

At cathode At anode 

fPb*+ + 2e- -> Pb fPb«+ + 2H .O - 2e~ -+ PbO, + 4H+ 
1504*- into solution \S0 4 *- into solution 
That is, the cathode grid acquires a filling of spongy leadmd the anode 
grid one of lead(IV) oxide. Passage of ions into solution (from equations 
above) in the proportion of 2S0 4 a ~ to 4H+ increases the concentration 
and density of the acid. At full charge, the E.M.F. is a little above 2 volts 
and the acid density is 1.25 gem - *. . .. «.,__, 

During discharge the cell yields electrical energy by the following 

At anode 
PbO, + 4H+ + 2e-— ► Pb'+ + 2H.O 
From solution, SO« ,_ . 
PbSO, deposits. 
Electrons available from lead at the cathode pass round the external 
circuit performing the electrical work required and are absorbed at the 
anode. Absorption of ions (4H+ : 2S0 4 2 -) from the electrolyte decreases 
the concentration and density of the acid. The E.M.F. falls to 2 volts 
soon after discharge begins and stays constant until it is almost complete, 
then railing to 1.8 volts. At this point, recharging is required. {For the 
Leclanchi cell, see p. 152) 

At cathode 
Pb->Pb ,+ + 2e~ 
From solution, SO/ - . 
PbSO, deposits. 

Chapter 15 

Water and Solution 

Water is essential to life 

WATER is of fundamental importance to ail kinds of plants and 
animals and therefore to man. It is of equal importance with the 
air we breathe in maintaining the vital processes necessary to life and 
growth, but since it is not everywhere available its provision has, 
from the earliest times, limited the setting up of villages and towns 
to the places where a water supply existed. Not only is water used 
all over the world in vast quantities for drinking purposes, but it is 
used in even greater quantities for washing, bleaching, dyeing, 
raising steam to drive engines, and as a solvent in industrial processes 
far too numerous to mention. It is the concern of the chemist to 
ensure that a supply of water is maintained which is suitable for all 
these purposes. 

If the water is too soft (see p. 201) it will attack the lead of the pipes 
in which it may be carried. If acids are present from decaying organic 
matter, sufficient lead may be dissolved by the water to cause harm 
to those who drink it. If the water supply for a town is too hard, 
because of the high percentage of dissolved solid matter which it 
contains, no firm will consider establishing a new industry there 
because of the enormous extra expense which they will incur in 
softening the water. 

Purification plants used in swimming baths, by the use of chlorine, 
keep the water comparatively free from the bacteria which carry 
many infectious diseases. The development of the high-pressure steam 
boiler, for the driving of machinery of all kinds, would have been 
impossible but for the solving of the problem of how to obtain water 
of a sufficiently high degree of purity for use in these boilers. The above 
are a few samples of how Chemistry is of use in the service of man. 

Occurrence of water 

Pure water does not exist in a natural state, but supplies of water 
are obtainable all over the world, varying in degrees of purity from 




rain water from clean districts (which contains 0.0005% of solid 
impurities) to sea-water, in which the impurities reach the compara- 
tively high proportion of 3.6%. (In certain lakes the proportion of 
solid matter is even higher.) 


A sample of fairly pure water can be made from rain water, tap 
water, or river water by the process of distillation. 

The impure water is placed in the distilling flask and is boiled (Fig. 
56). The steam comes off (together with gaseous impurities) whilst 


Fig. 56. 
Distillation of impure water. 

the solid impurities are left behind. The steam is condensed to water 
in the condenser. A pure liquid distils at a constant temperature 
which is its boiling-point at the prevailing pressure (for water 100°C 
at 760 mm). This behaviour is a recognised test for purity of a 
liquid. Tot fractional distillation, see p. 218. 

Pure water 

The preparation of perfectly pure water (or as near to perfectly 
pure water as its unusual solvent powers will allow) is a matter of 
much difficulty. It is prepared for conductivity experiments by as 
many as twenty successive distillations from a pure tin or platinum 



retort into a receiver which has been cleaned by having the purest 
water then obtainable kept in it for ten years! Potassium permangan- 
ate is added to the impure water in the earlier stages to oxidise 
organic impurities. 


Water is a clear colourless liquid with an insipid taste. It is usually 
recognised in the laboratory by its capacity to turn anhydrous 
copper sulphate (white) to a blue colour. 

CuSO« + 5Hj,0 -> CuSO.,.5H 2 
copper water hydratcd copper 
sulphate sulphate (blue) 

This test, of course, merely denotes the presence of water and not 
the absence of everything else except water; e.g., a dilute solution of 
sulphuric acid would turn anhydrous copper sulphate from white to 
blue. PURE water has the following properties: 

(a) It freezes at 0°C. 

(b) It boils at 100°C, when the barometer stands at 760 mm, and 
pure water will boil away completely with no change in temperature. 

(c) Its maximum density is 1 g/cm* at 4°C. 

(d) It is neutrarl to litmus. 

Water as the universal solvent 

There are few substances which do not dissolve in water to some 
extent. Even when you drink your morning glass of water, you are 

.Tap water 


Small paper- 

wwnnmtwirinnr &jiittc 

Appearance of clock-glass 
after evaporation 

Fia. 57. 
To show tap water contains dissolved solids. 



drinking a little of the glass as well. It is true you need not get 
alarmed, for the amount is very small indeed, but for certain ex- 
periments ordinary glass vessels cannot be used as containers for 
water because of this solvent effect. 

Tap water can easily be shown to contain a considerable quantity 
of both dissolved solids and gases by the following experiments. 

To show tap water contains dissolved solids 

Fill a large clean clock-glass with tap water and evaporate it down 
to dryness on a steam-bath as shown in Fig. 57. On holding the glass 
up to the light or against a sheet of white paper, you will observe a 
large number of concentric rings of solid matter left as the water 
gradually evaporated. 

To show that tap water contains dissolved gases 

Fill a flask with water and put in a few pieces of porous pot. Insert 
a two-holed rubber stopper to which are fitted a delivery-tube and 

Fio. 58. 
To show tap water contains dissolved gases. 

a short piece of glass tubing which can be closed by a rubber tube 
and clip and attached to the tap in the initial stages (the end of the 
delivery tube should not project beyond the surface of the stopper). 
By attaching the small piece of glass tubing to the tap by means of a 
rubber tube the whole apparatus (including the delivery-tube) is 
filled with water (Fig. 58). As the water is heated, bubbles of gas are 
seen to rise and these will collect and be carried over into the burette. 
Boil the water until no more gas is given off. The gas can be shown 
to differ from air in that its oxygen content is much higher, and the 
gas boiled out will rekindle a glowing splint. 




Woter level 

To determine the volume of oxygen in the air boiled out of water 

• Use in the last experiment a burette into which will fit the absorp- 
tion cup (G. Fowles) shown in Fig. 59. 

Fill the cup with crystals of 
pyrogallic acid and add water 
to fill up the air spaces. (See 
Fig. 59.) Have ready a piece 
of solid caustic soda (make 
certain that it will be able to 
enter the burette), insert it 
into the burette and follow it 
quickly with the absorption 
cup. Invert the tube several 
times and release the cup under 
water. Transfer the tube to a 
deep gas-jar, lower it until the 
levels are the same inside and 
out, and note the volume of 
gas absorbed. The percentage 
of oxygen in this air will be 
more than 30%. 

Value of these dissolved 
gases to fish life. The oxygen 
of the air dissolves in water 
to the extent of only 4 volumes 
of oxygen in 100 volumes of 
water (i.e., 1 dm* of water 
contains only a maximum of 
40 cm 3 [or 0.06 g] of oxygen). 
Although this amount is only very small it is of utmost impor- 
tance to fish life. The fish (with a few exceptions) rely on this 
oxygen for breathing, in just the same way as we rely on the air 
around us. 

Chemical value of the solvent properties of water 

Use is made in the laboratory of this exceptional property to bring 
into very close contact the particles of reacting substances. When the 
particles dissolve in water they have an opportunity for movement 
which they do not have in the solid state. Under these circumstances 
many reactions take place which do not take place if the reactants are 
solids (see ionisation), e.g., 

AgNO a (aq) + NaCl(aq) -»- AgCl(c) + NaNO a (aq) 

The above reaction will not take place if common salt and silver 

. — — — - -/— - -^■Caustic soda' stick 

Pyrogallol crystals 
covered with water 

Fig. 59. 
Absorption cup. 



nitrate are ground together in a mortar, because the average distance 
of the particles from one another is too great. 

In the above cases the water acts as a medium in which the action 
takes place. It does not as a rule react with ihe substances. The 
chemical actions of water, however, as an oxide and as a hydroxide 
producer are extensive. 

Action of water on metals 

r I Attack 
v ' a f water. 


Do not attack 
water or steam. 

By an examination of the electrochemical series it is easily shown 
that water attacks the metals to a degree varying with their position 

in the series. 
Potassium. Place a small 1 piece of potassium on water in a large 


Fio. 60. 
Action of potassium on water. 

dish (notice the silvery gleam of the unoxidised metal as it is cut 
with a knife). The potassium melts to a silvery ball, darts about the 
water and a gas is given off (hydrogen) which burns spontaneously 
with a violet flame (the colour is due to the burning of small quantities 
of potassium vapour) (Fig. 60). 

2K + 2H 2 -»- 2KOH + H 2 
potassium water potassium hydrogen 

1 Under no circumstances should the piece of potassium be larger than a small 
pea, as the action is very vigorous. Stand well back from the dish as the action 




If a piece of red litmus is placed in the dish it will turn blue because of 
the presence of potassium hydroxide, which is an alkali. 

Sodium. Perform the above experiment 
with sodium The sodium melts to a 
silvery ball, but does not burn unless it is 
restricted in movement. Effervescence 
occurs, a gas is liberated, and if a light is 
applied it burns with a yellow flame (the 
yellow colour is from the sodium). If the 
sodium is packed tightly into a sodium 
spoon (see Fig. 61) the gas can be collected 
and shown to be hydrogen. 

Fig. 61. 

Action of sodium 
on water. 

2Na + 2H a O 
sodium water 

* 2NaOH + H, 

sodium hydrogen 




Fig. 62. 

Action of calcium 
on water. 

Calcium. Drop a piece of calcium (a grey metal) into a dish of 

water and invert over it a boiling-tube full 

. H . of water. There is effervescence, and a gas 

V\ (hydrogen) is given off which explodes if 

§ jfr l mixed with air and a flame is applied (Fig. 

" | 62). The calcium dissolves, and if carbon 

dioxide is bubbled into a sample of this 
solution (blow into it down a glass tube), 
a milkiness due to a suspension of calcium 
carbonate in water is obtained. The solution 
is, in fact, lime-water, i.e., calcium 
hydroxide (slaked lime) dissolved in water. 

Ca + 2H s O -> Ca(OH), + H, 
calcium water calcium hydrogen 


Magnesium. This metal will not act appreciably on cold water but 
is attacked by boiling water. Steam vigorously attacks a piece of 
lighted magnesium ribbon. The demonstration of the action of 
magnesium on steam is beset with the practical difficulty that the 
temperature is so high that if the burning magnesium comes into 
contact with any glass, cracking immediately ensues. The following 
apparatus overcomes this difficulty in a simple way. 

A stout, very wide short-necked flask is partly filled with water and 
a piece of red litmus paper is dropped into the water and the water 
heated to boiling (Fig. 63). Meanwhile a piece of magnesium ribbon 
is attached to the piece of apparatus by pulling the cork away from 
the glass tube a little, and pushing the ribbon in parallel with the 
glass. Over the glass tube is placed a boiling-tube. When the water 





has boiled for a minute, remove 
the Bunsen and let bubbling 
cease. The ribbon is lighted and 
plunged into the steam, the cork 
fitting into the neck of the flask. 
The magnesium continues to 
burn, forming a white ash (mag- 
nesium oxide) which drops off 
into the water. When the action 
has ceased, remove the boiling- 
tube and apply it immediately 
to a flame, when there will be an 
explosion due to the combination 
of the hydrogen produced with 
oxygen of the air which is not 
completely displaced from the 

The piece of litmus paper will, 
after a time, turn blue. The mag- 
nesium oxide is only slightly sol- 
uble in water, but it will dissolve 
sufficiently to turn litmus blue, showing the presence of a hydroxide. 

Mg(c) + H.O(g) -»- MgO(c) + H 2 (g) 

(N.B. The above experiment shows that steam contains both 
oxygen and hydrogen.) 

Alternative method (see Fig. 64) 

Blow a small hole in a hard-glass test-tube by applying a small 
blowpipe flame until the tube is soft and blowing at the open end 

whilst the glass is being heated. 

Fig. 63. 
Action of burning magnesium on steam. 

Hydrogen bunts 
on meeting alr^ 

Put a coil of 15 cm magnesium 
ribbon in the tube, add 2 or 3 cm 3 
of water, insert the cork and clamp 
the apparatus by the cork at the 
angle shown (Fig. 64). Heat, gently 
at first, with a Bunsen burner, 
keeping the latter moving to 
maintain an atmosphere of steam. 
Finally heat strongly when the 
magnesium will burn and simul- 
taneously the liberated hydrogen 

will burn as it meets the outside air. 
Zinc does not attack hot or cold water. If zinc is heated to redness 

in a current of steam, hydrogen is formed. 


Bunsen mono" Co 
heat both water 
and magnesium 

Fig. 64. 


Zn + H,0 ->- ZnO + H, 


Iron does not attack water (rusting takes place only when air is 
present as well) but is readily attacked by excess of steam at red heat. 
This method is often used as a method of preparing hydrogen in 
quantity (see p. 265). 

3Fe + 4H 2 -> Fe,0 4 + 4H 2 
iron steam triferric hydrogen 

(red hot) tetroxide 

black oxide 
of iron 

The above reaction can be made to proceed in the reverse direction 
by passing excess of hydrogen over heated black oxide of iron (see 
p. 158). 

Action of water on non-metals 

Carbon attacks steam at a bright red heat, forming carbon mon- 
oxide and hydrogen. 

C(c) + H 2 0(g) -> CO(g) + H 2 (g) 

Chlorine acts on water to form hypochlorous acid. 

Cl s + H 2 0-»-HOCl + HC1 

chlorine water hypochlorous hydrochloric 

acid acid 

Action of water on oxides 

Potassium oxide is attacked by water with the formation of potas- 
sium hydroxide. 

K 2 + H 2 0-»-2KOH 
potassium water potassium 
oxide hydroxide 

Sodium oxide is similarly attacked by water with the formation of 
sodium hydroxide. 

Na.O + H 2 0-»-2NaOH 
sodium water sodium 

oxide hydroxide 

Calcium oxide (lime). Place a piece of quicklime in a dish and add 
water a few drops at a time. For a little while nothing is observed and 
then water-vapour is seen to come off, whilst a hissing sound as the 
water drops on indicates that the mass is becoming hot. It com- 
mences to expand and crack, and finally crumbles to a powder, 
slaked lime. 

CaO + H 2 -»• Ca(OH), 

calcium water calcium 

oxide hydroxide 




The above three hydroxides are soluble in water (slaked lime only 
slightly) and together with ammonium hydroxide form the common 

Sulphur dioxide reacts readily with water to form sulphurous acid. 

S0 8 + H 2 0->-H 2 SO, 


Similarly, other acidic oxides form acids with water, e.g., 

S0 3 + H 2 0->-H 2 SO, 

sulphur sulphuric 

trioxide acid 

2N0 2 + H 2 -> HNO, + HNO, 
nitrogen nitrous nitric 

dioxide acid acid 

C0 2 + H 2 -»- HgCO, 

carbon carbonic 

dioxide acid 

Action of water on chlorides. See Chlorides, p. 370. 

Action of water on certain metallic carbides 

Calcium and aluminium carbides react with water forming the 
hydroxide of the metal and hydrocarbons. 

CaC 2 + 2H 2 -> Ca(OH) a + C 2 H 2 


Al t C, + 12H 2 -»- 4Al(OH) 3 + 3CH« 


Composition of water 

The results of experiments already performed indicate clearly that 
water contains hydrogen and oxygen. If hydrogen is burnt in oxygen 
or exploded with it, water is produced (p. 269) and nothing else 
(p. 116). It remains to find thenumber of atoms of each element which 
is present in the molecule. 

Volume composition of steam. The experiment described on p. 116 
shows two volumes of steam to be formed from one volume of 
oxygen and two volumes of hydrogen. It follows that the formula for 
steam is H 2 0. 

Electrolysis of water. The above is the synthesis of water from its 
elements whilst electrolysis is the analysis. On electrolysis of water 
(see p. 145), it is found that the volume of hydrogen liberated is twice 
that of the oxygen liberated, thus confirming the above experiment. 

Gravimetric composition of water 

Hydrogen was at first taken to be the standard (H = 1) for the 
determination of Atomic Weights. Now the first step in a deter- 
mination of an atomic weight is usually a determination of the 





equivalent weight, and since many elements do not combine with or 
displace hydrogen readily, their equivalents have been determined by 
finding the weight of the element which combined with the equivalent 
of oxygen. Hence it is necessary to know with great accuracy the 
equivalent of oxygen. 

This is determined by passing purified hydrogen through several 
U-tubes of calcium chloride to dry it, and then through a tube con- 
taining pure dry copper oxide which is heated strongly. The latter 
tube and its contents are first weighed (Fig. 65). The water which 
passes over is collected in another set of calcium chloride tubes which 
are weighed before and after the experiment. 

CuO + H, — >- Cu + H 2 

copper hydrogen copper water 


The loss in weight of the tube containing the copper oxide gives the 
weight of oxygen, and the difference between the increase in weight 
of the absorption apparatus and the weight of oxygen gives the weight 
of hydrogen which has combined with the oxygen to form water. 

Pure dry 

Black Copper 
Oxide A 

J Absorption oF 
• water apparatus 

Calcium Chloride 
tube to prevent 
entry of 
moisture Into 

Calcium Chloride 
to absorb water 
Loss in weight gives weight of oxygen Gain in weight gives weight of 
which has combined with the hydrog en water formed 

Difference gives weight of hydrogen 

Fig. 65. 
Gravimetric composition of water. 

By an experiment involving the combination of the elements 
themselves it was found that 

1 g of hydrogen combined with 7.9396 g of oxygen 

This accurate figure was obtained by Morley (1890-5), who suc- 
ceeded in weighing the hydrogen, oxygen and also the water formed. 
The latter, of course, introduces a check on any error made. The diffi- 
culty of weighing a volume of hydrogen (because of the very small 
weight of even a large volume) was overcome by absorbing the 

hydrogen in the metal palladium. The hydrogen was expelled by 
heating the palladium to redness. 


There are many types of natural water found on the earth's surface. 

Rain water is the purest natural water, and if collected in a country 
district contains oxygen, nitrogen, carbon dioxide (dissolved as the 
rain drops pass through the atmosphere) and only a small amount 
of dissolved solids (0.0005%). 

River water, from which many domestic supplies are obtained, will 
obviously contain the same gaseous impurities and also any solids 
which the water has dissolved as it passed over the soil. The amount 
and kind of impurity will depend, therefore, on the type of soil over 
which the water runs. If the water runs over impervious material 
such as granite, the river water may be nearly as pure as rain water. 
In actual practice many springs and rivulets feed the large river from 
which a town's supply is obtained, and hence the impurities are often 
the same as those of spring water. 

Spring water is water which has made its way downwards thiough 
the soil and contains solid impurities. 

Sea-water is the reservoir into which all the impurities eventually 
go, and hence the solid content of sea-water is usually high (3.6%). 

The solids which are found in the natural waters are mainly the 
sulphates and bicarbonates (hydrogen carbonates) of calcium and 
magnesium together with smaller amounts of sodium chloride, 
silicates, nitrates, ammonium salts, as well as the gaseous impurities 
already mentioned as being present in rain water, i.e., oxygen, 
nitrogen and carbon dioxide. 

Of the solid impurities the most important are calcium sulphate 
and calcium bicarbonate (hydrogen carbonate), 

Calcium sulphate is present because many rocks and soils 
contain gypsum (CaS0 4 .2H g O) which is slightly soluble in 
water (1 : 500) and hence some of it dissolves in any water 
with which it comes in contact. 

Calcium hydrogen carbonate is present because water 
which contains carbon dioxide is capable of dissolving small 
quantities of limestone or chalk (which is found in large 
quantities in some soils and in small quantities in practically 
all soils). 

CaCO, + H t O + C0 2 -»• Ca(HC0 3 ), 
limestone water carbon calcium hydrogen 

(insoluble) dioxide carbonate 

(slightly soluble) 



It will be obvious that all the more soluble substances 
present on the surface of the earth were washed away in the 
course of past ages. 

Definition. A hard water is one which will not readily form a lather 
with soap. 

The nature and method of manufacture of soap 

Soap is the sodium salt of an organic 1 acid. One of the commonest 
soaps is the sodium salt of stearic acid, C 17 H M COi,H, and has the 
formula, C^HjaCOjNa, sodium stearate. It is important to notice 
that, for all its complex formula, soap is of exactly the same chemical 
nature as common salt, NaCl. Both are sodium salts of acids. The 
complex group, C^HajCOj— , of sodium stearate corresponds to the 
CI of sodium chloride. For convenience, the formula of sodium 
stearate is often written NaSt, the St being used as a substitute for the 
stearate group, C^HsjCOj — . 

Soap is manufactured by heating vegetable oils (such as palm oil 
or olive oil) or animal fats with caustic soda solution. The oils or 
fats are compounds formed from glycerine and certain complex 
organic acids, such as stearic acid, mentioned above. The caustic 
soda liberates glycerine from the fat and forms the sodium salt of 
the acid, which is the soap. (Caustic soda is sodium hydroxide.) 

Fat or oil + caustic soda — > soap + glycerine 
Soap is manufactured by steam-heating the fat and caustic soda solu- 
tion in large pans. Common salt is added later and assists in the 
separation of the soap which, when cool, sets as a hard cake on the 
surface of the liquid. It is removed and purified, dyes and perfumes 
being added to produce toilet soaps. 

The manufacturing process can be illustrated in the laboratory, 
using mutton fat or lard. Put the lard into an evaporating dish and 
add to it a solution of caustic soda to which methylated spirit has 
been added to quicken the action. Heat the dish on a steam-bath. 
When all the liquid has evaporated off, a yellowish solid will be left. 
It is impure soap. 

The soap cleanses by dissolving in the water, loosening the particles 
of dirt, and the whole (soap, water and dirt) can then be washed 

Cause of hardness 

Now if there happens to be a calcium compound dis- 
solved in the water, the soap is precipitated in the form of 

1 An organic compound is a compound of carbon. The chemistry of carbon 
is so complex that it is convenient to treat it separately from that of other 
elements as 'organic chemistry'. 


calcium stearate (which appears as a 'curd'), the latter being 
insoluble, e.g., 

2NaSt + CaS0 4 -*■ Na 2 S0 4 + CaSt 2 
sodium calcium sodium calcium 

stearate sulphate sulphate stearate 
(soluble) (soluble) (soluble) (insoluble) 

Until the whole of the calcium compound has been acted 
upon by the soap, none of the latter can form a lather. 
Thus, with a hard water, a large amount of soap is used to 
precipitate and remove the calcium, and only a small extra 
amount to cause a lather. 

In this way the valuable stearate group, which loosens the dirt, is 
lost completely, since it would be just as useful to try to wash with 
insoluble calcium stearate as any other substance insoluble in water, 
e.g., marble, or iron. 

This is not the only reason why hardness must be removed. Where 
water is used for boilers, certain of these solid substances (calcium 
sulphate and magnesium silicate being the worst offenders) are left 
behind as scale on the inside of the pipes of the boiler as the water is 
evaporated off. As this scale increases in thickness, the bore of the 
pipe becomes smaller, and the walls of which the pipes are made 
become heated to a higher temperature than is normal, causing the 
pipes to weaken and finally burst. 

Other detergents have been introduced recently to replace soap in 
domestic and laundry work. A typical detergent is made from a 
complex alkene, i.e., a hydrocarbon of the type, C„H 2 „, where n is 
between 12 and 20. The alkene is first sulphonated by concentrated 
sulphuric acid; the sulphonate is then converted to its sodium salt 
by sodium hydroxide solution. The sodium salt is the detergent. An 
example is the following. 

C„H„ + H 2 S0 4 -»- C 16 H 3I .S0 4 H 

C M H 3l .S0 4 H + NaOH -> C 15 H 31 .S0 4 Na + H,0 

This detergent resembles sodium stearate (soap) in possessing a long 
fat-soluble carbon chain, Ci 6 H 3I — , and a water-soluble end-group, 

— SO«Na. 

Detergents of this kind have two advantages over soap; they are 
more soluble in water and are not affected by hardness in water. This 
is because the calcium salt of the detergent is soluble in water whereas 
calcium stearate is insoluble. (See following sections.) 

Removal of hardness 

Distillation will remove all solid matter. This method is, as a rule, 
far too costly to be employed. 



Many high-pressure boilers supply steam to drive turbines, and the 
steam which comes from the turbines is condensed to water, which is 
actually distilled water. In some works, the water is fed back again 
into the boiler and small amounts of distilled water artificially pre- 
pared are added to make up the losses which are inevitable. 

Removal of calcium as calcium carbonate (chalk). Many of 
the methods of rendering a hard water soft have as their 
object the conversion of a soluble calcium salt into the in- 
soluble carbonate. In this way the calcium is removed, since 
the calcium carbonate, being insoluble, takes no further 
part in the reaction. An insoluble calcium salt cannot cause 

Temporary hardness 

Hardness which is due to the presence of calcium hydrogen car- 
bonate can be removed by heating the water to boiling for a few 
minutes. Heat decomposes the calcium hydrogen carbonate into 
calcium carbonate (chalk) and carbon dioxide is expelled. 

Ca(HC0 3 ) 3 -> CaC0 3 + H 2 + CO, 
calcium calcium water carbon 

hydrogen carbonate 
(slightly soluble) 



Because it can be removed merely by boiling, the name 'temporary' 
is given to this type of hardness. This method would be expensive on 
the large scale. 

Furring of kettles 

In a district where the water contains calcium hydrogen carbonate 
the insides of kettles become coated with a layer of calcium carbonate 
caused by the decomposition of the hydrogen carbonate according 
to the equation shown above. 

Stalagmites and stalactites 

These pillars of almost pure calcium carbonate are made by water 
containing dissolved calcium hydrogen carbonate dripping from the 
roof on to the floor of a cavern. Some of the calcium hydrogen car- 
bonate decomposes, giving off carbon dioxide into the atmosphere of 
the cave, and depositing calcium carbonate a little at a time on the 
roof and floor. This deposition causes a stalactite to grow downwards 
from the top of the cave and a stalagmite to grow upwards from the 
floor of the cave, until after a time the two meet. The growth 
varies very much from small fractions of a cm to 100 cm or more 
per year. 

If a sample of stalactite is available, its chemical nature can be 
demonstrated by the tests given below. 



(1) To a piece of the stalactite on a watch-glass, add a few drops 
of pure concentrated hydrochloric acid. Take up a little of the mix- 
ture on a platinum wire (sealed into the end of a glass tube to act 
as handle). Put the wire into the lower half of a Bunsen flame. A brick- 
red coloration shows calcium present. (At first, the yellow flame of 
sodium may interfere, because sodium compounds are very common 

(2) Put two or three small pieces of stalactite into a narrow test- 
tube and add dilute hydrochloric acid. Effervescence will occur. If the 
gas is passed into lime-water and the liquid is shaken, it will show a 
white turbidity. That is, carbon dioxide is evolved, proving the test 
material to be a carbonate. (A hydrogen carbonate also gives off 
carbon dioxide with this acid, but only two common hydrogen car- 
bonates exist as solid in ordinary conditions— NaHC0 3 and 
KHC0 3 .) 

C(V" + 2H + -»- H a O + CO a 
That is, the stalactite contains calcium carbonate, probably slightly 

In cross-section, stalactites usually show a ring structure similar 
to that of a tree trunk. This occurs because the summer rate of 
deposition of chalk is more rapid than the winter rate, 'emperature 
being lower in winter. 
Removal of temporary hardness by addition of slaked lime 

Temporary hardness can be removed by the addition of the cal- 
culated quantity of slaked lime (excess of lime would cause hardness 
on its own account). The amount of lime is calculated from a know- 
ledge of the hardness of the water and of the capacity of the reservoir 
(Clark's method). 

Ca(OH) 3 + Ca(HCO s ) 2 -»- 2CaC0 3 + 2H,0 
slaked lime calcium calcium water 

(slightly hydrogen carbonate 

soluble) carbonate (insoluble) 

(slightly soluble) 
A third method of removal of temporary hardness is to add sodium 
carbonate. (Permanent hardness is also removed at the same time.) 
Na,C0 3 + Ca(HC0 3 ),->- 2NaHC0 3 + CaCO s \ 

Permanent hardness 

This is due mainly to the presence of dissolved calcium sulphate, 
which cannot be decomposed by boiling, and hence the name 'per- 
manent' is given to this type of hardness. 

It is most easily removed by the addition of washing soda crystals 
to the water, when a precipitate of insoluble calcium carbonate is 
thrown down and thus prevents the calcium from interfering. 





(You should make yourself quite certain of the reason why soluble 
calcium sulphate can cause hardness, whereas insoluble calcium 
carbonate cannot. See p. 204.) 

CaSO, + Na,C0 3 

calcium sodium 

sulphate carbonate 

■*■ CaC0 3 + Na,SO. 




To make temporarily hard water 

Bubble carbon dioxide into lime-water in a flask for about 20 
minutes. The precipitate at first formed will have dissolved to form a 
temporarily hard water. 

Ca(OH), + CO,-)- CaCO, + H a O 
lime- carbon calcium water 

water dioxide carbonate 

CaC0 3 + H,0 + CO, ->■ Ca(HC0 3 ) a 

calcium water carbon calcium 

carbonate dioxide hydrogen 

To make permanently hard water. Add a little gypsum or anhydrous 
calcium sulphate to water in a flask. Shake and allow to stand, 
decant olTthe clear liquid; you will find it will be permanently hard 

To make a soap solution. Scrape shavings off a tablet of pure Castile 
soap. Weigh out about 6 g of shavings and add to about 100 cm 3 of 
methylated spirit in a beaker on a water bath. Warm and stir. When 
dissolved, transfer to a litre flask and add water to make up about 
one litre. Shake and allow to stand. 

N.B. Since the composition of soap is variable the above solution 
need not be made up with great accuracy, nor are the readings to be 
taken as a mathematical comparison of hardness. In these experi- 
ments the hardness of various waters is being compared. 

Fill a burette with the soap solution and place 25 cm 3 of distilled 
water by means of a pipette into a conical flask. Run in 1 cm 3 of the 
soap solution at a time and, between additions, cork up the flask and 
shake. When a lather is obtained, which persists unbroken for two 
minutes, the titration is completed. You will easily see the difference 
between a 'curd' (which is formed when soap solution is run into a 
hard water) and a lather, as the former breaks very quickly and the 
latter consists of many tiny bubbles which reflect the light from the 
windows. When you have run 30 cm 3 of soap solution into a hard 

water, you need not proceed further. The water is hard. Having 
obtained the result for distilled water perform the following ex- 

To show that boiling softens temporarily hard water. Take 25 cm 3 of 
the temporarily hard water made as indicated above and titrate 
against soap solution. Notice the curd which forms. 

CatHCOj), + 2NaSt-> CaSt, + 2NaHC0 3 

Take another 25 cm 3 and heat to boiling on a tripod and gauze. 
Notice that a milkiness appears (chalk). Cool. Titrate against soap 
solution again and note your result when a lather is obtained which 
remains unbroken for 2 minutes. 

Ca(HC0 3 ),-»- CaCO a + H,0 + CO, 

To show that washing soda crystals soften permanently hard water 

Titrate 25 cm 3 of the permanently hard water and record your 
result. Take another 25 cm 3 of the hard water, add a few crystals of 
washing soda, and shake until they dissolve. Again you will notice a 
turbidity due to a precipitate of chalk. 

Na,CO s + CaS0 4 -»• Na,SO« + CaCO, J 
Titrate again and notice your results. They should speak for them- 
selves. A typical set of results is given here. 

Amount of soap solution 
necessary to produce a 
lather to last unbroken 
for two minutes. 


25 cm* distilled water 

2nd reading 3.0 cm 
1st reading 1.0 cm 

25 cm' temporarily hard water 

2nd reading 33.0 cm'\ 
1st reading 3.0 cm*/ 

25 cm* temporarily hard water 
after boiling 
2nd reading 40.0 cm'\ 
1st reading 33.0 cm'/ 
25 cm' permanently hard water 

2nd reading 31.0 cm'\ 
1st reading 1 .0 cm*/ 

25 cm' permanently hard water to 
which washing soda was 
2nd reading 36.0 cm 3 \ 
1st reading 31.0 cm*/ 
1 Indicates no lather produced. No more soap solution was added— the water 
is extremely hard. 

2.0 cm' 

30.0 cm' 

7.0 cm' 

30.0 cm' 

5.0 cm' 



Pcnnut it method of softening water 

The above methods of softening water (i.e., boiling and adding 
'soda') are used mainly in the home for softening small amounts of 
water. In the treatment of larger supplies of water (but not so large as 
to be treated by the lime method) the permutit process is used. Many 
'water softeners' sold for domestic use work on this principle. 

Permutit is a complex substance (hydrated sodium aluminium 
silicates) but we can regard it as Na 2 Y (Y = Al„Si 2 O g .xH 2 0). When 
a dissolved calcium salt runs over it, ion-exchange occurs. 

(Na + ) 8 Y*- + Ca 2, S(V-- 

■Ca^Y*- + (Na+),S0 4 2 - 

The sodium permutit will finally become a calcium permutit, and it 
can be made fresh again by running concentrated brine over it and 
washing away the soluble calcium chloride formed. 

CaY + 2NaCl -> Na,Y + CaCl, 


Place a few crystals of copper(II) sulphate in a test-tube and heat 
gently. There will be a copious evolution of water-vapour (which will 
condense and run back and crack the tube if the mouth of the tube is 
not held lower than the bottom). The colour and shape of the crystals 
will disappear, and in place of blue crystals of hydrated copper sul- 
phate, you will observe a white powdery mass of anhydrous copper 

CuS0 4 .5H 2 — >- CuS0 4 + 5H,0 

If water is slowly added to the white, anhydrous powder, hissing is 
heard and steam is evolved, showing that heat is generated. A blue 
solid is left. This liberation of heat energy occurs because the an- 
hydrous copper sulphate is hydrated by the water and this is a 
chemical action, i.e., combination of H a O molecules with the ions 
of the salt. The heat generated is called heat of hydration. The 
reaction shown in the above equation is reversed. 

Water of crystallisation is necessary to the crystalline shape of 
some crystals, and is that definite amount of water with which the 
substance is associated on crystallising out from an aqueous solution. 
The crystals cannot form in these cases without the presence of water 
with which to form a loose compound. It is sometimes termed 'water 
of hydration'. 



These contain water of 

Sodium carbonate crystals 

(Na 2 CO a .10H 2 O) 
Sodium sulphate crystals 

(Na a SO 4 .10H 2 O) 
Copper sulphate crystals 

(CuS0 4 .5H 2 0) 
Iron(ll) sulphate crystals 

(FeS0 4 .7H 2 0) 

These do not contain water 
of crystallisation 

Sodium chloride 

Potassium permanganate 

KMn0 4 
Potassium nitrate 

KNO a 
Ammonium sulphate 



Many solids possess the property of dissolving in water. For in- 
stance, if we shake up a few small crystals of copper sulphate (crush 
them in a mortar first) with water in a test-tube, the water will turn 
blue and the crystals will finally disappear. 

The particles of copper sulphate in the solution must be very small, 
for we cannot see them with the naked eye, nor even with the most 
powerful microscope made. Furthermore, they will pass through the 
pores of a filter-paper with ease. The solution (as this mixture is 
called) of copper sulphate in water is uniform in blue colour and in 
composition (after stirring) and, in a sealed tube, remains so in- 
definitely. In general, to produce such a solution, the ordered crystal 
lattice of the solid (the solute) built up from positive and negative 
ions (see NaCl, p. 371) or from molecules (see I t , p. 75), must first 
be broken down so that its particles can diffuse between the mole- 
cules of the liquid (the solvent) and move among them. Since the 
particles in the solid lattice exercise electrical attraction on each 
other, energy must be supplied to separate them. This energy is 
taken from the heat energy of the liquid; consequently, there is 
usually a fall of temperature as a solid dissolves in water. If, how- 
ever, some form of chemical action occurs at the same time, e.g., 
hydration of ions, this action may liberate enough heat to cause an 
overall rise in temperature, e.g., when anhydrous copper sulphate 
is dissolved in water. 

Liquids other than water may act as solvents, e.g., chloroform, 
carbon tetrachloride and carbon disulphide. These liquids, and many 
others, are covalent compounds and dissolve a great number of 
covalent solids, e.g., chloroform dissolves iodine, carbon disulphide 
dissolves sulphur. Ionic compounds, e.g., Na + Cl - , do not usually 
dissolve in covalent solvents but a great many are soluble in water, 
in which their ions can dissociate. 



To obtain the solute from a solution 

Into a porcelain evaporating dish pour a solution of common salt 
in water, and heat the contents on a tripod and gauze. Water-vapour 
will be seen to come off and finally there will be left a white solid, 
common salt. 

This method can be used to purify rock salt, since earthy impurities 
and dirt will not dissolve and can be removed by filtration. 

To obtain the solvent from a solution 

If we could catch the water-vapour coming off in the previous 
experiment and cool it, we could recover the water free from common 
salt, since the solid does not come off with the vapour. The method 
of catching the vapour and cooling it is called condensation and the 
apparatus employed is a distilling flask, condenser and receiver (see 
p. 191). The condenser consists of a long tube surrounded by a 
wider tube through which water can flow, to cool the vapours and 
convert them into a liquid again. The whole process of boiling off 
the liquid as vapour and then cooling it again so that it is received as 
liquid in another vessel is called distillation. The product obtained 
here is called 'distilled water'. It is free from dissolved solids but may 
contain gases in solution. It has a 'flat' taste which makes it un- 
palatable. Distilled water is used in electric accumulators because 
the solids present in ordinary tap-water are injurious to the plates. 

Chemical solution 

This is a term often employed to indicate the apparent solution of 
a solute in a solvent, together with chemical action. For example, zinc 
appears to dissolve in dilute sulphuric acid and neither zinc nor dilute 
sulphuric acid can be recovered by evaporation or distillation, since 
the solid residue on evaporation would be zinc sulphate. Actually the 
processes of chemical action and solvent action follow one another. 
The zinc attacks the acid to form zinc sulphate, which then dissolves 
in the water present. 


The above properties of a solution help us to differentiate between 
a solution and a suspension. A solid is said to be in suspension in a 
liquid when small particles of it are contained in the liquid but are not 
dissolved in it. 

If the mixture is left undisturbed the solid particles will 
slowly settle to the bottom of the containing vessel, leaving 
the pure liquid above them. 



Muddy water is a typical suspension. The mud would settle after a 
time if left undisturbed, leaving a brown residue on the bottom of the 
containing vessel and clear water above. The particles of mud would 
be retained by a filter-paper whilst the water (and any solids in 
solution) would pass through. 

Saturated solution 

If we add half a gram of sodium chloride to 100 g of water in a 
beaker the salt will dissolve. We could go on for a time adding salt 
half a gram at a time and, by stirring vigorously after each addi- 
tion, bring about solution of the sodium chloride, but with increasing 
difficulty. Finally, there would come a time when no more sodium 
chloride would dissolve at that particular temperature, and, no matter 
how long we left it or how vigorously we stirred, no more sodium 
chloride would dissolve. The solution is then said to be saturated with 
sodium chloride at the particular temperature. 
Definition. A saturated solution of a solute at a particular temperature 
is one which contains as much solute as it can dissolve at that tem- 
perature, in the presence of the crystals of the solute. 

The concentration of a saturated solution varies with the solute, 
the solvent and also with the temperature. Thus sulphur is almost 
insoluble in water yet readily dissolves in carbon disulphide and a 
rise in temperature will cause more to dissolve. 

This is generally true for solids; for example, nitre is at least seven 
times more soluble in water at 80°C than it is in water at 10°C. 

Determination of solubility 

To give a quantitative meaning to solubility, it is necessary to fix 
the amount of the solvent and to state the temperature under con- 
sideration. The amount of solvent is usually fixed at 100 g. 
Definition. The solubility of a solute in a solvent at a particular tem- 
perature is the number of grams of the solute necessary to saturate 100 g 
of the solvent at that temperature. 

It denotes a limit, that is, the maximum amount which can normally 
be held in solution. Solubility is also sometimes expressed in grams 
of solute per litre of solution at a given temperature. 

To determine the solubility of potassium nitrate in water at the tem- 
perature of the laboratory 

This determination must be carried out in two stages. It is first 
necessary to prepare a saturated potassium nitrate solution at labora- 
tory temperature and then to find the proportions of potassium 
nitrate and water in it. 

To make the saturated solution. The rate of solution of a solid in 



cold water is generally so slow that it is almost impossible to obtain a 
saturated solution of it in a reasonable time by merely shaking the 
solid with the water. The quicker and more certain way is to crystal- 
lise from a warm solution by cooling. 

Half fill a boiling-tube with water and dissolve in it some potassium 
nitrate. Warm and shake well. Pour off a small sample into a test-tube 
and cool it under the tap. If no crystals appear, return the sample 
to the boiling-tube and add more nitre. Test another sample 
and continue in this way till a sample gives crystals. Then cool the 
whole solution. When the crystals have separated and the solution is 
quite cold, take the temperature of it. Then filter it through a dry 
filter-paper and funnel into a dry receiver to avoid diluting it. The 
filtrate is a saturated solution of potassium nitrate at the observed 

To obtain the solubility of potassium nitrate, using this solution. 
Weigh a clean dry dish and add some of the saturated solution to it. 
Weigh again. Once having weighed be careful not to lose any portion 
of the solution. Place the dish on a steam bath (p. 192), and evaporate 
until the potassium nitrate is left quite dry. (The dish on a gauze may 
be warmed very gently over the Bunsen flame for a few minutes to 
complete the removal of water.) Allow the dish to cool and weigh it. 
Calculate the weight in grams of the potassium nitrate which would 
have dissolved in 100 g of water as in the following calculation. 

Alternative method of evaporation to dryness. The following method 
is quicker than the one suggested above with little loss of accuracy if 
carefully performed. Weigh a dish (75 mm diameter is suitable) with 
a clock-glass to fit over it. Weigh again, having added some saturated 
solution of potassium nitrate and replaced the clock-glass. Evaporate 
to dryness over a medium flame increasing the size of the flame 
towards the end. Do not heat the solid sufficiently strongly to decom- 
pose it. Allow to cool and weigh. 


Weight of dish = 14.32 g (a) 

Weight of dish and solution = 35.70 g (o) 

Weight of dish and potassium nitrate = 18.60 g (c) 

Temperature of saturated solution = 15°C 


17.10 g (b — c) of water dissolve 4.28 g (c — a) of potassium nitrate 

4 28 
so, 100 g of water dissolve -^-— x 100 g of potassium nitrate 

= 25.0g 



That is, solubility of potassium nitrate is 25.0 g per 100 g of water at 

Generally, increase in temperature increases the solubility of a 
solute in water. This is because most solutes dissolve in water with 
absorption of heat. So, if a solution and solid crystals of the solute 
are in equilibrium at a certain temperature and the mixture is then 
heated, the system will alter so as to tend to lower the temperature 
again (Le Chatelier's Principle, p. 156). That is, more crystals will 
dissolve because their dissolution absorbs heat. 

By finding the solubilities of a solute at varying temperatures a 
graph can be plotted to show how the solubility alters with increase 
of temperature, with many interesting results. This is called a solu- 
bility curve of the solute. 

To determine the solubility of potassium nitrate at S0°C 

This is the method employed to determine the solubility at any 
temperature above laboratory temperature. At these higher tempera- 
tures, the rate of solution of the potassium nitrate is greatly increased, 
and a saturated solution may be made directly. 




Fio. 66. 
Solubility of potassium nitrate at 50 'C 



Crush some crystals of potassium nitrate in a mortar, place some 
of them in a boiling-tube and add a little water (to make the tube 
about half-full). Put the boiling-tube in a beaker of water and warm 
the latter up to a temperature of about 55°C (Fig. 66). Whilst warm- 
ing the solution keep adding potassium nitrate crystals to the boiling- 
tube, and stir all the time. Add potassium nitrate until some remains 
undissolved at the bottom of the tube. Remove the flame when the 
temperature, as read by the thermometer, reaches 55°C, and allow 
the apparatus to cool, stirring all the while and always maintaining 
some undissolved potassium nitrate at the bottom of the tube. Just 
before the temperature falls to 50°C, remove the stirrer, allow the 
solid potassium nitrate to settle, and put the dry thermometer into 
the potassium nitrate solution. When the temperature is exactly 50°C, 
rapidly decant a little of the saturated solution into a weighed dish, 
leaving all solid potassium nitrate behind. Weigh the dish again and 
evaporate to dryness as in the previous experiment. Calculate the 
weight of potassium nitrate dissolved in 100 g of water at 50°C in a 
similar way. 

By repeating the experiment at varying temperatures several values 
can be obtained from which a curve can be plotted. On p. 215 is a 
table of values obtained in this way. Plot the graph on squared paper. 
The following figures were actually obtained by a middle school 
form. (The figures in brackets are the accepted accurate values.) 

Temperature ll°C 

g KNO, per 
100 g water 

/23.6 g 

| L20 g) 



(85 g) 


102 g 
(106 g) 

Alternative method (using any substance which does not exhibit 

Weigh 4.5 g of potassium chlorate into a boiling-tube and run in 
10 cm 3 (g) of water from a burette. Warm until dissolved, remove 
from the flame, insert a thermometer and allow to cool, stirring with 
the thermometer. Note the temperature at which crystals appear. 
This will be the temperature at which the solubility is 45 g. Add a 
further 10 cm 3 of water and repeat the experiment. Continue the 
addition, determining the temperature at which the solution is just 
saturated until 60 cm 3 of water have been added. Construct the graph. 

Graph of solubilities 

Look at the accompanying graph of the solubilities of a few com- 
mon salts. (Fig. 67.) Answer the following questions. 


Solubility k 
in grams 
solute per 

100 g jo 





zo 30 40 bo si 

Fio. 67. 
Solubility curves of a few common substances. 

(a) For which salt does the solubility increase most rapidly with rise 
in temperature? 

(b) Given a mixture of equal weights of potassium chlorate and 
potassium chloride, how would you obtain some pure chlorate? 

(c) For which salt is there a decrease in solubility with increase in 


Fill a 150 X 25 mm boiling-tube to a depth of about 25 mm with 
water, then fill it up with crystals of 'hypo', sodium thiosulphate, 
NajS 2 3 .5H 2 0. Heat the tube gently and shake until the crystals 
are all dissolved. Holding the boiling-tube quite still, cool the solution 


No crystals 




's spreading 

-downwards after 

m GO 


Fig. 68. 



under the tap. Even when the solution is quite cold, no crystals separ- 
ate (Fig. 68, i). Now select a single pin-head size crystal of 'hypo' and 
drop it into the solution. At once, white crystals separate, growing 
slowly downwards through the solution and starting from the 
added crystal as centre (Fig. 68, ii). The contents of the boiling-tube 
become almost a solid mass of 'hypo' crystals and only a very little 
solution can be poured off. Note the rise of temperature as the 
crystals separate. 

After the crystals have separated, the solution left must be still 
saturated with 'hypo'. Before they separated, it must have been 'more 
than saturated', or 'super-saturated', as it is called. 
Definition. A solution is said to be super-saturated when it contains in 
solution more of the solute than it can hold at that temperature in the 
presence of the crystals of the solute. 

It is important to notice that 'super-saturated' solutions are in an 
unstable condition. They can only be obtained to any marked extent: 

(a) From a few compounds, for example: Glauber's salt 
Na.SO^.lOHjO, 'hypo', Na 2 S 3 3 .5H t O. 

(b) By excluding all dust. The dust particles might act as centres 
of crystallisation, 

(c) By cooling the solution slowly. 'Hypo' is exceptional in giving 
a super-saturated solution when cooled quickly. 

(</) By avoiding all shaking or disturbance of the solution. 
A solution cannot be super-saturated if in contact with crystals of 
the solute. 

Effect of dissolved solids on the boiling-point of a solution 

Put about 200 cm 3 of water into a distilling flask (500 cm 3 ) fitted 
with a sensitive thermometer, which should read to tenths of a degree. 
Add a few pieces of porous pot, clamp the flask at a suitable height 
and heat the water to boiling-point and notice at what temperature 
the water boils. Add a weighed amount of sodium chloride, say 5 g 
and again determine the boiling-point. Repeat this experiment, adding 
exactly 5 g of salt each time and noting the boiling-point after each 

Specimen readings 

Temperature of boiling water 99.7 C C 

With 5 g of NaCI added 100.2°C 

„ 10 g 100.7°C 

„ 15 g „ „ „ 101 2°C 

From these figures, elevation of boiling-point is proportional to 
concentration (if solutions remain dilute and at constant pressure). 



Effect of dissolved substances on the freezing-point of solutions 

Water, if pure, freezes at 0°C, but if it contains some dissolved 
substance, it is found that the solution no longer freezes at 0°C, but 
at a lower temperature. It is also found that equal small additional 
quantities of the dissolved solid cause equal depressions of the freez- 
ing-point. Hence, the depression of the freezing-point of a solvent is 
proportional to the mass of a substance dissolved in a given mass of 
the solvent for dilute solutions at constant pressure. 

Determination of melting-point 

The melting-point of a solid (between, say, room temperature and 
150 o C) can be determined as in Fig. 69. A thin glass melting-point 



cone, sulphuric 
acid or glycerol- 



Test material 

Fio. 69. 
Determination of melting-point. 

tube about 4 cm long (which can be drawn out from ordinary glass 
tubing) is touched into a Bunsen flame to seal one end. When it has 
cooled, a little of the finely powdered solid, e.g., naphthalene or 
benzoic acid, is scooped into the open end and tapped down to the 



bottom of the tube to make a 2-3 mm layer. The tube will adhere to 
the thermometer by the surface tension effect of the liquid in the 
flask and the test material should be as close as possible to the ther- 
mometer bulb. The flask is gently heated and the melting-point tube 
is carefully watched. At a certain point, the solid will be seen to melt 
and the corresponding temperature (melting-point) is immediately 
read. A pure solid will melt completely over a very narrow tempera- 
ture range (less than 0.5°C). Impurity depresses the melting-point and 
usually causes gradual softening instead of sharp liquefaction. To 
melt sharply at its recognised melting-point is good evidence of purity 
for a known solid. 

Fractional distillation 

If two liquids have reasonably different boiling-points, e.g., water 
100°C, ethanol (alcohol) 78°C at standard pressure, they can be 
separated (though not always completely) by fractional distillation. 
Various forms of fractionating column can be used (Fig. 70). Their 
general purpose is to provide surfaces, e.g., flat discs, on which 



Fio. 70. 
Fractional distillation. 



ascending vapour can condense. This produces a succession of liquid 
films in which, as the column is ascended, there is an increasing con- 
centration of the more volatile liquid (of lower boiling-point). Ascend- 
ing vapour comes into approximate equilibrium with these liquid 
films. Consequently, the vapour which leaves the column and passes 
to the condenser is far richer in the more volatile constituent than 
the original mixture. 

For example, a mixture of equal volumes of 'methylated spirit* and 
water will not burn in air. If it is distilled as in Fig. 70, and few cm 3 
of the early distillate are put into an evaporating dish, the liquid will 
burn readily on the application of a light. That is, the proportion of 
alcohol is significantly increased. An efficient industrial still can 
deliver alcohol of about 96% concentration from a very dilute alcohol 
wash produced by fermentation, but this mixture has a lower boiling- 
point than any other water-alcohol mixture and distillation cannot 
purify it further. 

An industrial fractionating column (as used, e.g., in distillation of 
crude petroleum) will be cylindrical, made of steel and up to, per- 
haps, 60-70 m high. Hot vapour ascends through the middle of the 
column. Trays at regular intervals round the walls collect layers of 
condensed liquid through which ascending vapour is forced to pass 
by baffles. The column delivers fractions of distillate from various 
points, with decreasing boiling-points higher up the column. See also 
p. 324. 

Determination of water of crystallisation 

Barium chloride crystals, BaCI 2 .xH 2 0, provide a suitable case. A 
clean dry crucible is weighed with lid. Two or three grams of barium 
chloride crystals are added and the whole is weighed. The crucible, 
with lid, is then heated on a pipe-clay triangle on a tripod, gently 
at first and later strongly, to drive off water of crystallisation. The 
whole is allowed to cool in a desiccator (to exclude moisture) and 
weighed. It is then heated again, cooled as before and re-weighed. 
This is repeated until a constant weight is reached which shows that 
all water has been expelled. 


Wt. of crucible and lid a g 

Wt. of crucible and lid and barium chloride crystals b g 
Wt. of crucible and lid and barium chloride anhydrous c g 

Wt. of water expelled (b — c) g 

Wt. of barium chloride anhydrous (c — a) g 
xUfi (b - c) 


BaCl, (c - a)' 


and, inserting molecular weights, 

18s _ {b-c) _ {b - c) 208 

208 "(T^) and x ~(F=7) x 1z 


1. Explain the difference between (a) temporary, and (b) permanent 
hardness in water. Give the names of one substance causing temporary 
and of two substances causing permanent hardness in water. Explain how 
permanent hardness may be removed, giving equations for the chemical 
changes involved. How would you test for the presence of a chloride in 
tap-water? (N.U.J.B.) 

2. What are solubility curves? Of what use are they? What experiments 
would you make in order to construct a solubility curve for potassium 
chlorate? (O. and C.) 

3. What are the substances which cause (a) temporary; (6) permanent 
'hardness' in water? 

Explain the effect of adding the following substances to 'hard' water: 

(a) soap; (b) lime-water; (c) washing soda. 

What is the 'fur' deposited in the kettle? Briefly explain how stalactites 
and stalagmites are formed in caves. (B.) 

4. What are the conditions, and what are the products, for the reaction 
of the following substances with water or steam: (a) iron; (6) calcium; 
(c) charcoal; (a") chlorine; (e) calcium carbide? (D.) 

5. Imagine that you wish to prove experimentally to someone that 18 
g of water contain 2 g of hydrogen and 16 g of oxygen. Give a labelled 
sketch of the apparatus you would use; indicate the two chief precautions 
you would adopt to ensure an accurate result; show how you would use 
the data you obtain to prove the above statement. (N.U.J.B.) 

6. Ordinary tap-water always contains some air in solution. Describe in 
detail how you would collect a quantity of this air from tap-water. How 
could you find the proportion by volume of oxygen in such air? Explain 
why the composition of this dissolved air will be different from that of 
ordinary air. (N.UJ.B.) 

7. State and explain what happens when carbon dioxide gas is passed for 
a long time through lime-water. State and explain what happens when soap 
solution is shaken with the final product (a) before it has been boiled- 

(b) after it has been boiled. (N.U.J.B.) 

8. What is meant by the terms 'saturated solution' and 'super-saturated 
solution'? Describe exactly how you would proceed to determine the 
solubility of potassium nitrate in water at 15°C. (N.U.J.B.) 

9. What experiments would you make to find out whether (a) a given 
sample of water is hard or soft; (6) a gas-jar contains nitrous oxide or a 
mixture of nitrogen and oxygen; (c) a given solid is sodium sulphide or 
sodium sulphate? (O. and C.) 

10. Choose either (a) or (b), but not both: 

(a) Write an essay on 'combustion', giving a brief historical intro- 



(b) Define the term solubility. What are the effects of temperature 
and pressure on the solubility of (a) gases; (b) solids, in water? 
Give the requisite practical details for constructing a solubility 
curve of a salt, e.g., nitre. State two of the uses of a solubility 
curve. (B.) 

1 1. A colourless liquid (X) is either pure water, or water containing some 
dissolved solid. Describe carefully how to discover which it is by obser- 
vations of the temperature at which it (a) boils; (6) freezes. Sketch the 
apparatus you would use in each case. Assume that a supply of pure water 
is provided. (N.U.J.B.) 

12. What do you observe when carbon dioxide is passed for a long time 
through lime-water? Give equations representing the reactions which take 

What happens when the final solution obtained is (a) treated with soap 
solution; (b) boiled; (c) treated with a solution of sodium carbonate. 

Describe the 'permutit' process for softening water and state how the 
permutit is restored (or revivified). (C.W.B.) 

Chapter 16 

Acids, Bases and Salts 


ACIDS have a variety of properties (as stated later), but all these 
properties are derived from a single type of behaviour. This be- 
haviour is the production of the hydrogen ion, H + , by acids when 
dissolved in water. Consequently, an acid can be defined in the 
following way : 

Definition. An acid is a compound which, when dissolved in water, 
produces hydrogen ions, II , as the only positive ion. 

Examples are : 

Hydrochloric acid 1 
Sulphuric acid 
Nitric acid 

HC1 ^ H + + Cl- 
H,S0 4 ^ 2H+ + S0 4 «- 
HNO s ^ H + + N0 3 - 

The simple hydrogen ion (proton) is hydrated by water to form the 
hydroxonium ion, H 3 + , but this complication is often ignored in 
the elementary consideration of acidity. 

H+ + H,0 ^ H„0 + 

If, as in the above three dilute acids, the ionisation is almost complete, 
the acid is said to be a strong acid. If the ionisation is only slight, the 
acid is said to be weak. For example, bench acetic acid is about 
0.4% ionised, that is, only four out of every thousand molecules of 
the acid pass into the ionised state. 

CH 3 COOH ^ H + -f- CH 8 COO" 
Per thousand mol. 996 4 4 

Some properties of acids 

Taste. Most dilute acids have a sour taste. This is true of the three 
common mineral acids, sulphuric, hydrochloric and nitric, and of 
many others. The sour taste of many unripe fruits, lemons and sour 
milk is caused by the acids in them. 

1 For HCI in toluene, sec p. 240. 



Action on litmus. Most acids turn blue litmus to red. Some of the 
weaker acids, however, for example carbonic acid, are so feebly 
acidic that they can only turn the litmus to claret colour. 

Corrosive action. The man in the street connects the term acid 
with the idea of a corrosive, 'burning' liquid. This is because two of 
the commonest acids — sulphuric acid (oil of vitriol) and nitric acid 
(aqua fortis)— are actually corrosive liquids. Acids are not, however, 
generally corrosive and most of them are solids. 

Action with metals. Metals which are much more electropositive 
than hydrogen (p. 134), react with dilute hydrochloric or dilute 
sulphuric acid (or both) to liberate hydrogen. The metals, e.g., Zn, 
Mg, Al (with HCI only) and Fe, supply electrons which are taken 
up by H + ions of the acid, as: 

Zn -»- Zn 2+ + 2e~ 
2H + + 2«r-»-H 2 

A metal such as lead, which is only slightly more electropositive than 
hydrogen, reacts with neither of these dilute acids, but will react in a 
similar way with hot, concentrated hydrochloric acid. 

Nitric acid is too strongly oxidising to allow the liberation of 
hydrogen. Oxides of nitrogen are obtained. For a further account of 
this, see p. 437. 

Action with carbonates. Almost all acids (only the very weakest 
are exceptions) liberate carbon dioxide from a carbonate. The carbon- 
ate supplies the ion, C0 3 *-, which, with H + from the acid, gives the 
reaction : 

2H + + CO s *--> H.O + CO, 

Action with bases and alkalis. This is very important but must be 
postponed (to p. 224) till the nature of bases and alkalis has been 

Methods of preparation of acids 

(1) By the reaction between an acid anhydride (the acidic oxide of 
a non-metal) and water. 

Examples of this method are the preparation of sulphurous acid 
and metaphosphoric acid by the action of sulphur dioxide and 
phosphorus pentoxide with cold water. 

SO s + H t O ^ H 2 S0 3 ; P 4 O 10 + 2H a O -»- 4HP0 8 

(2) By displacing a weaker or more volatile acid from its salt by a 
stronger or less volatile acid. 

For example: (a) displacement of the more volatile hydrogen 
chloride from metallic chloride by the less volatile concentrated 
sulphuric acid. 

NaCl + H 2 SQ 4 -»- NaHSO« + HCI 



(b) displacement of the weaker boric acid from borax by sulphuric 

Na,B 4 7 + H s SO< + 5H t O 


Na,SO« + 4H 8 B0 3 
(boric acid) 

(3) By precipitating an insoluble sulphide from a metallic salt by 
hydrogen sulphide. 

Pb(C a H,O t ) t + H 1 S- 

(lead acetate) 

PbS + 2C a H«0, 
(acetic acid) 


A little earlier, an acid was characterised as a producer of H + in 
aqueous solution. To correspond with this, a base may be defined in 
the following way: 

Definition. A base is a substance which can combine with hydrogen 
ion, H+. 

A number of bases are mentioned below and their combination 
with H + is shown. The reactions are reversible. If, as the equations 
are written below, the equilibrium lies strongly to the left, i.e., few 
molecules arc formed, the base is said to be weak; if the equilibrium 
lies strongly to the right, i.e., few ions remain, the base is said to be 

of Base 

very weak 

very weak 

very weak 


very strong 

Chloride ion 
Nitrate ion 
Sulphate ion 
Acetate ion 
Hydroxyl ion 


CI" + H + 
NO a - + H l 
SCV" + 2H + 
CH3COO- + H+ 
OH-+ H* 



HN0 3 

H 2 S0 4 


H 8 

It is obvious that a base exists corresponding to each acid. The 
corresponding pairs can be called conjugates and, if the acid is weak, 
the base is strong, and vice versa. 

At the chemical level for which this book is intended, the most 
important of all the bases is the hydroxyl ion, OH". It is a very 
strong base indeed, and when it combines with hydrogen ion, the 
water formed is almost completely unionised, i.e., from this point of 
view, water acts as a very weak acid. Metallic oxides and hydroxides 
are usually electrovalent compounds, containing the ion, 2 ~ 
(oxides), or OH- (hydroxides), in addition to the metallic ion. The 
ion, O*-, can enter the equilibrium: O*- + H g O ^ 20H~ with 
water. Consequently, metallic oxides and hydroxides of this kind 
can produce the base, OH~, and are known as basic oxides or 



hydroxides. By reacting with an acid, i.e., with H + , they produce 
molecules of water. 

H + + OH- ^ H,0 

At the same time, the metallic ion from the oxide and the negative 
ion from the acid remain, and together constitute a salt. Examples 
of this are: 

Na+OH- + H+Cl--*- Na'Cl- + H 2 
Ca 5+ (OH-) 2 + 2(H+N0 3 -)-v Ca st (NO a -) 2 + 2H,0 

From these considerations, the following definitions can be derived: 
Definition. A basic oxide (or hydroxide) is a metallic oxide (or hydrox- 
ide) which contains ions, O 2- (or OH"), and will react with an acid 
to form a salt and water only. 

It is important to realise the importance of the word only in these 
definitions. If it was omitted, certain compounds, which are quite 
different from basic metallic oxides and hydroxides, would be in- 
cluded under the definition of base. Thus, lead(IV) oxide reacts with 
hydrochloric acid to produce lead(II) chloride (a salt) and water, 
but the word only excludes it from the class of bases because chlorine 
is also produced. 

PbO a + 4HCI -> PbCl, + 2H 2 + Cl» 

Compare this with the equations above. Lead(IV) oxide is clearly 
not a base. 

The reactions between basic oxides, or hydroxides, and acids are 
very important and are called neutralisations. Since the metallic ions 
and anions from the acid do not change (see the two equations above), 
the essential reaction of a neutralisation is always the formation of 
unionised molecules of water from hydroxide and hydrogen ions 
(or hydroxonium ions). 

H + + OH-^H g O or H s O + + OH~ ^ 2H 2 

For practical purposes, these reactions are complete from left to 
right. This leads to the following definition. 

Definition. Neutralisation is the formation of molecules of water from 
hydroxide ions, OH - , and hydroxonium ions, H 3 '. A salt is formed 
at the same time. 

The following are examples of neutralisations. Notice that the 
metallic ions and the negative ions from the acids remain to produce 
the salts. 


Basic hydroxide 




+ K+OH- 

— ► 

K + C1- 

+ 2H a O 

H 3 + N0 3 - 

+ Na + OH- 

— >■ 

Na + N0 3 - 

+ 2H,0 



SinceboththesereactionsreducetoH s O + + OH-— *-2H 8 0,theenergy 
change accompanying both should be the same. This is, in fact, so. 
If aqueous acidic solutions are made up containing one mole of 
HCI, HNO s and $H 2 S0 4 in 1 dm 3 and alkaline solutions containing 
one mole of NaOH and KOH in 1 dm 3 (i.e., strong acids and alkalis), 
any neutralisation between these solutions liberates the same amount 
of heat energy, viz., 57.3 kJ (13.7 kcal). This was surprising in an 
earlier 'molecular' context because all the compounds appeared to 
have separate identities. Ionic ideas, however, require all strong 
acids and alkalis (and the salts they produce) to be completely ionised 
in dilute solution. Consequently, the only significant change between 
them in neutralisation is: 

H,0+ + OH - -»- 2H 2 0; A// = -57.3 kJ 


A definition of basic hydroxide has been given above. If such a 
hydroxide is soluble to a considerable extent in water, it is known as 
an alkali. This gives the definition: 
Definition. An alkali is a basic hydroxide which is soluble in water. 

Only a few alkalis are known. The common ones are sodium 
hydroxide (caustic soda), NaOH, potassium hydroxide (caustic 
potash), KOH, calcium hydroxide (slaked lime), Ca(OH) 2 , and 
ammonia solution. This property of solubility in water is the only 
difference between the select little group of alkalis and the basic 
hydroxides generally, but it is a very important difference. It puts the 
alkalis at our service in hundreds of reactions in solution and for a 
great many purposes for which the insoluble basic hydroxides are 
quite useless. It must be clearly realised that the alkalis have all the 
properties of basic metallic hydroxides in general, but possess also 
the property of dissolving to a substantial extent in water. Slaked 
lime is the least soluble of the common alkalis (about 0.15 g in 100 g 
of water at room temperature). 

Properties of alkalis 

In addition to the very important property of neutralising acids, 
already dealt with above, alkalis also have a bitter taste and turn red 
litmus blue. Further, the two caustic alkalis, NaOH and KOH, have 
a powerful corrosive action on the skin and should be treated with 
care. Just as a non-volatile acid displaces a volatile acid from its 
salts, a non-volatile alkali displaces the volatile ammonia from 
ammonium salts. If an ammonium salt is warmed with an alkali (in 
the presence of water), ammonia gas is liberated. The essential reac- 
tion is: 

NH 4 < +OH--*NH 3 + H 2 

SALTS 227 

Expressed in molecular form, two examples of this are: 

NaOH + NH 4 CI ->■ NaCl + H a O + NH 3 


Ca(OH) 2 + (NH 4 ) 2 S0 4 ->- CaS0 4 + 2H a O + 2NH 3 


We have encountered the term salt on p. 225, applied to a product 
of the process of neutralisation. A salt was seen to consist, from this 
point of view, of an aggregation of metallic ions (positively charged) 
and acidic ions (negatively charged). Consequently, a salt is derived 
from the acid to which it corresponds by replacing the H ' of the acid 
by an equivalent number of metallic ions. Examples are: 

(H*),S0 4 2 - 
H + NO a - 
(H+) 2 W" 

Na + Cl _ sodium chloride 
Cu 24 S0 4 2 - copper sulphate 
K + N0 3 _ potassium nitrate 
Ca z+ CO a 2 - calcium carbonate 

Notice that the salt is electrically neutral by a balancing of the 
oppositely charged ions. 

Normal and acid salts 

When an acid can produce more than one hydrogen ion per mole- 
cule, it is possible for the replacement of H+ by metallic ion to occur 
in stages. Thus, one of the H + ions of sulphuric acid, (H + ) 4 S0 4 2 -, 
may be replaced by the ion, Na + , to give sodium hydrogen sulphate, 
Na + HS0 4 -, after which a second, similar replacement may yield 
sodium sulphate, (Na + ) 2 S0 4 2 -. Salts like Na + HS0 4 " partake of 
the nature of a salt because they contain a metallic ion and a negative 
ion derived from an acid ; they partake of the nature of an acid because 
the negative ion is capable of further ionisation to yield H+, as: 
HSO t _ ^ H+ + S0 4 2_ . Having this dual nature, they are called 
acid salts. Salts like sodium sulphate, in which all the H + of the acid 
has been replaced by metallic ion, are known as normal salts. 
Definition. A salt is a compound consisting of positive metallic ions 
and negative ions derived from an acid; if the negative ions are capable 
of further ionisation to yield H* the salt is an acid salt, but if not, the 
salt is a normal salt. 

(H + ),S0 4 «- 

Acid salt 
Na+HS0 4 - 

(sodium hydrogen 

Normal salt 

(Na + ) e S0 4 *- 

(sodium sulphate) 


Acid Acid salt Normal salt 

(H+) a CO,»- Na^HCOr (Na+) t CO,«- 

(sodium hydrogen (sodium 

(H + ) S S 2 - 


(sodium hydrogen 

(Na + ),S 2 - 
(sodium sulphide) 

A method of preparing the normal and acid sodium salt of 
sulphuric acid is given on p. 412, of the normal and acid sulphides 
of sodium on p. 400 and of sodium carbonate and hydrogen car- 
bonate on p. 312. 

Basicity of an acid 

We have seen that it is characteristic of an acid to yield H* in 
aqueous solution. The number of H+ ions produced per molecule of 
the acid is called its basicity. 

Definition. The basicity of an acid is the number of hydrogen ions, 
H*, which can be produced by one molecule of the acid. 

Acid Basicity 

HC1 ^H + +C1- Monobasic 

H 2 S0 4 ^ 2H + + S0 4 «- Dibasic 

H 3 P0 4 ^ 3H+ + P0 4 3 ~ Tribasic 

Notice that the basicity of an acid is not necessarily the number of 
hydrogen atoms contained in one molecule of it. For example, acetic 
acid, CgH 4 2 , though containing four hydrogen atoms per molecule, 
is only niowobasic. Three of the four hydrogen atoms are so com- 
bined as to be incapable of ionising. 

QH 4 8 c* H+ + C.HA- 

Basic salts 

Basic oxides or hydroxides (p. 225) contain the ions, O 2- or OH - , 
respectively. Salts in which these ions are retained, together with 
metallic ions and the negative ions of acids, are known as basic salts. 

For example : 

Basic hydroxide Basic salt Normal salt 

Zn^(OH-), Zn* + OH-Cl- Zn 2+ (Cl-) t 

Mg 2+ (OH-)„ Mg 2+ OH-Cl- Mg a+ (Cl-) 2 

The basic salts quoted in the table are known as basic zinc (or 
magnesium) chloride. White lead is a well-known basic salt and is the 
basic carbonate of lead, 3Pb 2T 20H-2C0 3 8 -. Notice that the ionic 
charges balance so that the salt as a whole is electrically neutral. 



Practical details of the more important methods of making salts 
are given in the following pages. 


Rules of solubility of salts in water 


1. All common salts of sodium, 
potassium and ammonium. 

2. All common nitrates of metals. 

3. All common chlorides except 

4. All common sulphates except 


silver chloride, mercurous chlor- 
ide and lead chloride. 

barium sulphate and lead sul- 
phate. (Calcium sulphate spar- 
ingly soluble.) 

5. All common carbonates except 
those of sodium, potassium 
and ammonium (see Rule 1). 

Several general methods are available for preparing salts. The 
method chosen for preparing any particular salt depends largely on 
whether it is soluble in water or not. It is necessary, therefore, for 
you to become quite familiar with the simple rules of solubility 
above. Knowing the solubility of the salt, you will then be able to 
decide at once what type of method to use. 

Soluble salts are usually prepared by methods which involve 
crystallisation. Insoluble salts are usually prepared by methods which 
involve precipitation. 

Preparation of soluble salts by the action of an acid upon a metal 

This is usually carried out using a dilute acid and a metal. The salt 
formed then passes into solution in the water of the dilute acid and 
can be obtained, with the necessary precautions to give purity, by 
crystallisation. This method is suitable for preparing soluble salts. 

Preparation of a sample of crystalline zinc sulphate (white vitriol) 
from metallic zinc. Half-fill a beaker with bench dilute sulphuric 
acid and add granulated zinc. Effervescence occurs; if it is slow, add 
a little copper sulphate to quicken the action and warm gently. The 
gas is hydrogen (test: explosion on applying a lighted taper). If the 
action ceases because of the disappearance of zinc, add more zinc to 
make sure that the acid is not left in considerable excess. The reason 
for this is that later we shall evaporate the solution and any excess 



acid would then tend to become concentrated. When the effervescence 
slows down and there is still plenty of zinc left, filter to remove in- 
soluble impurities such as excess zinc and particles of carbon which 
were an impurity in it. 

The colourless solution will contain zinc sulphate together with a 
little sulphuric acid. There will probably be too much water present 
to allow the crystals to separate, so we must remove some of it. Place 
the solution in an evaporating dish and heat. At intervals, pour a 
little of the solution into a test-tube and cool it under the tap, shaking 
well. After evaporating for a time, you will see small crystals in one 
of the cooled samples. This shows that the solution will give crystals 
when cool, for the small sample is typical of the whole. Pour the 
solution into a beaker and allow it to cool and crystallise. In this 
case the crystals will be colourless needles. 

Filter off the crystals, wash them two or three times with a small 
quantity of cold distilled water and place them on a porous plate or 
between filter-papers to dry. The washing with distilled water is to 
remove the surface solution from the crystals and replace it with pure 
water which, as it dries off, will not deposit impurities as would the 
solution. The porous plate or filter-papers are used to absorb water 
from the surfaces of the crystals. To purify still further, dissolve the 
crystals in a little very hot water and repeat the crystallisation process. 
The impurities will then be carried off dissolved in the filtrate and a 
smaller quantity of purer crystals will be left. 

Zn + H 2 S0 4 -> ZnSO« + H 2 
(or Zn + 2H + -v Zn*+ + H 2 ) 
ZnS0 4 + 7H g O ->- ZnS0 4 .7H 2 

Other salts may be similarly prepared, for example, iron(II) sul- 
phate and magnesium sulphate. 

Preparation of iron(II) sulphate crystals (green vitriol). Use iron 
filings or wire and excess dilute sulphuric acid. The solution and 

crystals are green. 

Fe + H 2 S0 4 -»-FcS0 4 -f-H 2 

(or Fe + 2H+ -*■ Fe*+ + H 2 ) 

FeS0 4 + 7H s O-> FeS0 4 .7H 2 

(iron(II) sulphate crystals) 

Preparation of magnesium sulphate crystals (Epsom salt). Use 
magnesium rod and dilute sulphuric acid. The solution and the 
crystals are colourless. 

Mg + H 2 S0 4 -»- MgS0 4 + H 2 
(or Mg + 2H+ -»- Mg 2+ + H 2 ) 
MgS0 4 + 7H 2 -> MgS0 4 .7H 2 

(magnesium sulphate 



Occasionally this method is applied using a metal and a concen- 
trated acid. The only case of this you will encounter, though it is an 
important one, is 

Preparation of a sample of blue vitriol (crystalline copper sulphate) 
from metallic copper. Note especially that this preparation cannot be 
carried out like the three above, because dilute sulphuric acid will not 
act upon copper. 

Into a beaker put some concentrated sulphuric acid and coppet 
turnings (care!). Put the beaker on a tripod and gauze in the fume- 
cupboard and warm gently. After a time effervescence will begin, the 
gas evolved being sulphur dioxide, which must not be allowed to 
escape into the laboratory as it is injurious. The action will probably 
become vigorous, and the flame should then be removed. After 
effervescence has ceased there will be left a dark brown mass which 
may contain the following substances: 

1. Solid copper sulphate as a precipitate. (This is copper(II) 

2. Solid dark brown copper(I) sulphide, formed in small amount, 
by a side reaction. 

3. Excess copper. 

4. Excess concentrated sulphuric acid. 

Our problem is to prepare from this mixture a sample of pure blue 
vitriol. We cannot work in the presence of a large amount of con- 
centrated sulphuric acid, so the first step must be to pour off as much 
liquid as possible. The liquid is simply a waste product. We are now 
left with a solid containing copper sulphate with impurities 2 and 3, 
as given above. They are both insoluble in water, so, to remove them, 
add a considerable quantity of water and heat gently to boiling. 
Filter, leaving the two impurities, copper and copper(I) sulphide, on 
the filter-paper, and obtaining, as filtrate, blue copper sulphate 
solution. From this, crystals can be obtained as previously described. 
Cu + 2H 2 S0 4 ->- CuS0 4 + 2*1*0 + S0 2 
CuS0 4 + 5H 2 0-> CuS0 4 .5H,0 

This process of heating copper with concentrated sulphuric acid is 
the best laboratory method of making sulphur dioxide, and copper 
sulphate can be prepared from the residue left in the flask after carry- 
ing out this preparation (p. 403). 

Nitrates of certain metals can be prepared by acting upon the metals 
with dilute or concentrated nitric acid and crystallising as described 
(p. 230). The nitrates of common heavy metals, except lead nitrate, 
are, however, very soluble in water and deliquescent. This makes it 
a matter of greater difficulty to prepare their crystals. 

3Pb + 8HN0 8 -► 3Pb(N0 3 ) 2 -f 4H 2 + 2NO 



3Cu + 8HNO a -> 3Cu(NO a ) 8 + 4H 2 + 2N0 
copper coppcr(II) nitrate 

Chlorides of heavy metals are generally prepared in the anhydrous 
state by heating the metal in a current of dry chlorine or hydrogen 
chloride (see p. 372). 

The reason for this is that many of them crystallise with water of 
crystallisation, and, if an attempt is made to drive off this water by 
heating, they hydrolyse to basic salts (see p. 228). 

Thus: ZnCl 2 .H 2 -^ Zn(OH)Cl 4- HC1 

Preparation of salts by double decomposition 

In a double decomposition reaction, we usually begin with two 
soluble compounds and use them to prepare one soluble and one in- 
soluble compound (p. 160). Of these, the one which is wanted is 
usually the insoluble compound for it can be easily separated by 
filtering. That is, insoluble salts are prepared by double decom- 
position. For a more complete discussion of double decomposition 
see p. 160. In what follows, all lead salts are lead(II). 

To prepare a sample of lead sulphate. Here we must begin with a 
soluble lead salt to provide the lead ions, and a soluble sulphate to 
provide the sulphate ions of the lead sulphate. Suitable compounds 
will be lead nitrate and dilute sulphuric acid (a soluble sulphate). 

One-third fill a beaker with lead nitrate solution. Heat it and add 
dilute sulphuric acid, stirring the mixture. There will be an immediate 
white precipitate of lead sulphate. Heat to boiling, 1 then filter. The 
lead sulphate is left on the filter-paper and the colourless filtrate con- 
tains dilute nitric acid, the other product of the double decomposition 
reaction. Wash the precipitate on the filter-paper several times with 
hot distilled water to remove soluble impurities. To be sure that the 
process is complete, test the washings for sulphate (p. 416) and con- 
tinue till the test gives a negative result. Allow the precipitate to dry 
on the filter-paper or on a porous plate. It will become a white 

Pb(N0 3 ) 2 + H 2 S0 4 -»- PbSO« + 2HN0 3 

A solution of any soluble sulphate could have been used for this 
preparation; for example: 

Pb(N0 3 ) 2 4 Na 2 S0 4 -> PbS0 4 + 2NaN0 3 
lead sodium lead sodium 

nitrate sulphate sulphate nitrate 

ionically: Pb 2+ (aq) 4- S0 4 2 "(aq) -> PbS0 4 (c) 

1 The chief reason for this is that a boiling solution filters more rapidly than a 
cold one. 



The sodium nitrate, like the nitric acid above, would be removed in 


Other insoluble salts which can be prepared by double decom- 
position are: 

BaCl 2 4- H 2 S0 4 ->- BaS0 4 + 2HC1 

Ba*+(aq) + S0 4 *"(aq) -»- BaS0 4 (c) 
Pb(N0 3 ) 2 + 2NaCl -> PbCl 2 + 2NaNO s 
Pb s+ (aq) + 2Cl~(aq) -» PbC! 2 (c) 

->- CaC0 3 4 2NaCl 

Barium sulphate 

Lead chloride 1 

Calcium carbonate CaCl 2 4- Na 2 CO s 


ionically: Ca 2+ (aq) 4- C0 3 ~(aq) -*- CaCO s (c) 

The carbonates of other heavy metals can be prepared like calcium 
carbonate but, in some cases, the method gives a basic carbonate. 
This does not matter much, however, because the method is chiefly 
used as an intermediate stage in preparing the oxide of a metal from 
one of its soluble salts (p. 243), and in this case it does not matter 
whether a true or a basic carbonate is precipitated. 

In the cases of zinc carbonate, copper carbonate and lead car- 
bonate, a purer product is obtained if sodium hydrogen carbonate is 
used instead of sodium carbonate. 

Preparation of salts by neutralisation 

Neutralisation is an action between a base and an acid to produce a 
salt and water only (p. 225). The actual method of application of this 
process depends on whether the base in question is soluble in water, 
that is, is an alkali (p. 226) or not. 

Preparation of a salt from an alkali (a soluble base) 

Salts of sodium, potassium and ammonium can be prepared 
by this method from caustic soda (sodium hydroxide), 
Caustic potash (potassium hydroxide) and ammonia (ammo- 
nium hydroxide) respectively, using the appropriate acid. 

To prepare sodium sulphate from caustic soda. Into a conical flask, 
put 50 cm 8 of bench caustic soda solution and add to it enough 
litmus solution to give a pale blue liquid. From a burette, add dilute 
sulphuric acid until the solution is purple in colour; that is, it con- 
tains no excess of either acid (which would make the litmus red), or 

1 Lead chloride must be washed with cold distilled water, as it is appreciably 
soluble in hot water. 



alkali (which would make the litmus blue). The litmus is here the 
indicator (p. 239). The solution now contains sodium sulphate solu- 
tion together with litmus. To remove the litmus, add a little animal 
charcoal on the end of a spatula, boil the mixture, and filter. The 
animal charcoal will be left on the filter-paper together with the 
litmus which it has absorbed, and a colourless solution of sodium 
sulphate will be left as filtrate. From this, crystals can be obtained as 
described above for zinc sulphate (p. 230). 

2NaOH + H 2 S0 4 -»- Na 2 S0 4 + 2H 2 
Na 2 S0 4 + 10H a O-> Na 2 SO 4 .10H 2 O 

To obtain sodium chloride or nitrate, use caustic soda solution with 
the appropriate acid. 

NaOH + HC1 ->• NaCl + H 2 

sodium dilute sodium 

hydroxide hydro- chloride 

chloric acid 

NaOH + HNO s - 
nitric acid 

NaN0 3 + H 2 



To obtain the common potassium salts, use caustic potash solution 
with the appropriate acid. 

KOH + HCI -»- K.C1 + H 2 
potassium potassium 

hydroxide chloride 

2KOH + H 2 S0 4 -»- K 2 SO« + 2H t O 


KOH + HN0 3 ->- KNO a + H t O 


Ammonium salts can be similarly prepared from ammonia. 
NH 4 OH + HCI -*• NH 4 C1 + H a O 

2NH 4 OH + H 2 S0 4 

NH 4 OH + HN0 3 - 


(NH 4 ) 2 S0 4 + 2H 2 

NH 4 NO a + H t O 


The above equations indicate the substances to be used (all in 
solution) in the preparation of the salt by the method described for 
sodium sulphate. 



preparation of a soluble salt from an insoluble base 

In this case the base, being insoluble in water, is not an alkali, so 
the method given above cannot be applied. In what follows, all 
copper and lead compounds are copper(II) and lead(II) compounds. 

To prepare copper sulphate crystals from the insoluble base, copper 
oxide Heat some dilute sulphuric acid in a beaker and add to it, a 
little at a time, some black copper oxide. Stir gently. A blue solution 
of copper sulphate will be formed. 

CuO + H 2 S0 4 ->- CuS0 4 + H,0 
It is advisable not to leave an excess of acid (for reason see prepara- 
tion of zinc sulphate from zinc), so continue the addition of the 
copper oxide until a permanent black precipitate of this material is 
left showing that no more acid is available to act with it. Filter off 
the'precipitate, leaving a clear blue filtrate, copper sulphate solution. 
From it, obtain crystals as described for zinc sulphate. 
CuSO« + 5H 2 -*■ CuS0 4 .5H 2 

Many salts of metals can be similarly prepared, using either the 
oxide or the hydroxide of the metal with the appropriate acid; for 

Zinc sulphate ZnO + H 2 S0 4 -> ZnS0 4 + H 2 

H 2 S0 4 -*ZnS0 4 + 2H B 

Lead nitrate 

Zn(OH) 2 


PbO + 2HN0 3 - 

Pb(OH), + 2HN0 3 


Pb(N0 3 ) 2 + H 2 

Pb(N0 3 ) 2 -I- 2H 2 

Preparation of salts by the action of an acid on the carbonate of a 

The carbonate of any metal will react with the mineral 
acids to give the corresponding salt of the metal, water and 
carbon dioxide. 

For example : ZnC0 3 + H 2 S0 4 -> ZnSO« + H»0 + CO s 

zinc z'nc 

carbonate sulphate 

CaCO a + 2HC1 ->- CaCl, + H 2 + CO, 

calcium calcium 

carbonate chloride 



PbC0 3 + 2HN0 3 - 


► Pb(NO s ), + H s O + CO.. 
Icad(ll) nitrate 

The only limitation to this rule is that the action is unsatisfactory 
and incomplete if the carbonate is insoluble in water and, by its 
action on the acid, produces a salt which is also insoluble. In this 
case, the salt which is formed precipitates on the unchanged carbon- 
ate and stops the action. If, for example, dilute sulphuric acid is 
added to marble, there is rapid effervescence for a few seconds, but 
the action quickly stops. 

CaCOj + H,S0 4 -»• CaS0 4 + H,0 + CO, 

The very slightly soluble calcium sulphate has precipitated on the 
marble, stopping the action. With this limitation, this method is of 
general application. Lead compounds below are lead(II) compounds. 
Preparation of lead nitrate crystals from lead carbonate. Half-fill a 
beaker with dilute nitric acid, warm it gently on a tripod and gauze 
and add lead carbonate at intervals. There will be effervescence with 
evolution of carbon dioxide. (Test: lime-water turned milky.) Con- 
tinue the addition of the carbonate until a slight permanent precipi- 
tate shows that no acid is left, then filter to remove insoluble materials. 
The colourless solution contains lead nitrate, and crystals can be 
obtained from it as described for zinc sulphate (p. 230). 

PbC0 8 + 2HNO a ->- Pb(N0 3 ), + H,0 + CO, 

Note. When an insoluble salt has to be prepared from a 
compound which is also insoluble in water, it is not ad- 
visable to try to convert one into the other by a single pro- 

For example: To convert lead oxide, PbO (insoluble in water) 
into lead sulphate (also insoluble in water), it is not advisable 
to attempt the change by the action of dilute sulphuric acid on the 
oxide. The reason is that the insoluble lead sulphate will precipi- 
tate, as fast as it is formed, round the particles of oxide, and it is 
very difficult to make sure that all the oxide has reacted. 

In such cases, it is better to prepare first a soluble compound and 
then to precipitate the required insoluble salt by double decomposi- 
tion. In the example above, it is better to prepare first a solution of 
lead nitrate by dissolving the oxide in dilute nitric acid. 

PbO + 2HN0 3 -» Pb(N0 3 ), + H,0 

Then, after filtering off any undissolved material, the required lead 
sulphate can be precipitated by adding dilute sulphuric acid. 

Pb(NO s ), + H,S0 4 -»■ PbS0 4 + 2HNO, 

pH scale 237 

Similarly, to convert chalk to calcium sulphate, do not attempt the 
conversion by using dilute sulphuric acid directly with the chalk. First 
dissolve the chalk in dilute nitric acid. 

CaC0 3 + 2HN0 3 ->- Ca(N0 3 ) 2 + H,0 + CO, 
Then filter, and, to the hot filtrate, add a concentrated solution of 
sodium sulphate. 

Ca(NO s ), + Na.,S0 4 -> CaS0 4 + 2NaNO a 
Filter the precipitated calcium sulphate (or lead sulphate above) wash 
it with hot distilled water and allow it to dry. 

Preparation of salts by direct combination of two elements 

Certain binary salts can be prepared by direct combination of their 
two elements; for example, 

2Fe + 3Cl,->- 2FeCl 3 (see p. 362) 
iron(lH) chloride 
Fe + S -*- FeS (see p. 7) 
iron(ll) sulphide 

pH scale. This scale is a convenient means of expressing acidity and 
alkalinity in aqueous liquids. The term, pH, denotes hydrogen ion 
index, the p being derived from European use of terms such as punkt 
or point for the English index. The pH scale can be used with the mini- 
mum information contained in the next paragraph but students are 
advised to read the later material which explains more adequately the 
source of the scale and the significance of the numbers in it. 

The number 7 on the pH scale represents the condition of pure 
water in relation to acidity and alkalinity, i.e., the condition of exact 
neutrality. Numbers less than 7, i.e., pH 6, 5, 4, etc., indicate acidity 
increasing as the numbers decrease; numbers greater than 7, i.e., pH 8, 
9, 10, etc., indicate alkalinity increasing as the numbers increase. 
pH 1 2 3 4 5 6 7 8 9 10 11 12 13 

« t ; — ; — > 

increasing Neutral increasing 

acidity alkalinity 

Notice particularly that, below 7, falling pH values indicate in- 
creasing acidity; above 7, rising pH values indicate increasing alkal- 
inity. For the mathematical explanation of this, see later. In a colour- 
less liquid, a reasonably accurate value of pH can be obtained by 
adding a universal indicator (in quantity advised by the supplier) or by 
spotting the liquid on to universal indicator paper. In both cases, a 
colour will appear from which the pH of the liquid can be decided, 
either by verbal description or from a series of sample tubes provided 
by the supplier for colour matching. 






Ionic state of pure water. Hydrogen ion index, pH 

Pure water is very slightly ionised and is, therefore, a very weak 

H a O^H + + OH- or 2H s O ^ H s O + + OH~ 

(The first of these equations is still in common use for simplicity 
though it is known that the hydrogen ion (proton) is hydrated in 
aqueous solution to give the hydroxonium ion, H s O + .) From con- 
ductivity measurements, it is known that, in mole/dm s , the con- 
centration of the two ions in water is lO -7 , so that their product is 
given by: 

[H + ][OH-] = 10-' x 10-' = 10-" = K w (at 25°C) 
where K K is called the ionic product of water. This ionic product is 
maintained constant in all circumstances in water at 25°C. When the 
concentrations of H+ and OH~ are equal, the liquid is exactly 
neutral and both ion concentrations are 10~ 7 mole/dm 3 . If the liquid 
is made acidic, H + concentration is raised above 10~ 7 mole/dm 3 . 
Then, to maintain the value of K„, OH~ concentration must be 
reduced correspondingly below 10- 7 mole/dm 3 . For example, in 
0.01 M strong monobasic acid, [H+] = 10-*, so that 

10-°- x [OH-] = 10-" and [OH"] = 10"" 
If the liquid is made alkaline as, say, 0.001 M sodium hydroxide 
solution, [OH-] = 10" 3 , so, to maintain the value of K u [H,+] must 
decrease to 10 -u . This gives the following situation: 

Aqueous liquid H ' cone. 





above [10- 7 ] 
below [10- 7 ] 

OH - cone. 

[10- 7 ] 

below [10- 7 ] 
above [10- 7 ] 

As a convenient means of expressing these situations, the hydrogen 
ion index, pH, has been introduced. This index is derived in the fol- 
lowing way. 

If in an aqueous liquid, concentration of H+ is 10 - " 
mole/dm 3 , then, for that liquid, pH = x. 

From this it follows (above table) that neutrality is indicated by 
pH = 7. Notice that pH is a logarithmic index from which the 
negative sign has been omitted. Consequently, the greater the num- 
erical value of pH the lower is the concentration of H + and the less 
the acidity. 

It would be possible to use a hydroxyl ion index, pOH, to indicate 
alkalinity but, in fact, such a usage can be avoided. This is because, 
in all cases, 

[H + ][OH-] = 10-" = K v 
so that pH + pOH = 14 and pOH = 14 - pH 
This constant relation enables pH to record alkalinity (as well as 
acidity) as in the following case. 0.01 M potassium hydroxide solution 
contains 10" a mole/dm 3 of OH". This would give pOH = 2 and 
require pH = 14 - 2 = 12. That is, pH indicates alkalinity by 
values higher than 7. 





below 7 

above 7 

Universal indicator. A universal indicator is a mixture of indi- 
cators, often sold ready-made as a solution, which can indicate pH 
values, usually over a range of about 3-11, by successive changes of 
colour. A recognisable colour change takes place for every change of 
one unit of pH and, over certain ranges, only half a unit. 

A given solution is tested for pH by adding the universal indicator 
to it in the proportion stated by the supplier, usually two drops of 
indicator solution lo 10 cm 3 of test solution, with shaking. The 
colour which develops shows the pH of the solution either by recog- 
nition from a verbal description given on the indicator label or by 
matching to sealed sample tubes provided by the supplier for pH 
3, 4, 5, etc., up to 11. This method is suitable for colourless liquids 
only (for obvious reasons). Knowledge of pH values is important 
for many industries, e.g., for control of soil condition in farming and 
horticulture, for control of water supplies and sewage disposal. 

An ordinary single indicator usually requires a change of about 2 
units of pH to bring about its complete colour change. For example, 
litmus is fully red at pH 6, purple at pH 7 and fully blue at pH 8. 
The three common indicators for acid-alkali titration are the fol- 

Acidic side ' Neutral Alkaline side 

Methyl orange 






colourless ' colourless ' pink 

-2 pH units increase- 



Hydrogen chloride in toluene 

The solvent, toluene, C 4 H B .CH 3 , readily dissolves hydrogen 
chloride in room conditions, but (in contrast to aqueous hydrogen 
chloride solution) the liquid is a non-conductor of electricity and has 
negligible acidic behaviour, e.g., no reaction with metals (Zn, Mg, Fe) 
and only very slight reaction with carbonates. These differences are 
present because toluene is not a proton (H + ) acceptor and the dis- 
solved hydrogen chloride remains overwhelmingly in the molecular 
state as HCI; the solution is, therefore, electrically non-conducting 
for lack of ions. 

Water, however, is a proton acceptor and, in dilute aqueous solu- 
tion, hydrogen chloride can be taken as fully ionised. 

HCI + H a O -► H 8 0+ + Cl- 

These ions allow ready electrical conductance through the liquid 
and the hydroxonium ion liberates hydrogen with the more electro- 
positive metals and carbon dioxide with carbonates. 

Zn + 2H 3 0+ -»- Zn* + + 2H,0 + H, 

2H s O+ + CO,*--*- 3H,0 + CO, 

These reactions are not given in toluene for lack of H a O + . If 
ammonia (gas) is passed into it, a solution of hydrogen chloride in 
toluene precipitates ammonium chloride (white solid). 

NH 3 + HCl->-NH 4 +Cl- 


1, Describe fully, giving complete experimental details, how you would 
prepare the following. Consider in each case the solubility of the required 
compound and decide first which of the methods given in the last chapter 
is the best for the given case. In some of them it is advisable to use two 
steps (see Note, p. 236). 

(a) Copper(II) sulphate crystals starting from malachite, basic copper(II) 
carbonate, CuCO,.Cu(OH),. 

(b) lron(ll) sulphate crystals starting from iron wire. 

(c) Nitre (potassium nitrate). 

(d) Calcium sulphate starting from marble (calcium carbonate). 

(e) Sal-ammoniac (ammonium chloride). 
(/)Lead(II) chloride starting from lead(II) carbonate. 
(g) Zinc sulphate crystals starting from zinc oxide. 
(h) Sodium chloride, common salt. 

2. Describe fully, with complete experimental details, how you would 
prepare (a) copper(II) sulphate crystals from a mixture of charcoal and 
black copper oxide; (.b) zinc sulphate crystals from the commercial product, 
'zinc dust', which contains zinc and zinc oxide; (c) a pure sample of 
barium sulphate and zinc carbonate, normal or basic, from a mixture of 
zinc oxide and barium carbonate. 



3 Describe one laboratory method of preparing specimens of each of 
the 'following: (a) copper from copper(II) sulphate; (b) solid sodium 
hvdrogen sulphate from sodium hydroxide; (c) potassium chlorate from 
potassium hydroxide; (d) anhydrous iron(III) chloride from iron. (B.) 

4. Give four general methods of preparing acids, illustrating each 
method by two examples. . . 

Describe experiments by which you could determine the basicity of 
sulphuric acid. (L.) 

5. Describe, giving experimental details, how you would prepare 
crystalline specimens of (a) potassium nitrate starting from potassium 
hydroxide; (6) lead(II) nitrate starting from red lead ; (c) sodium hydrogen 
sulphate starting from sodium hydroxide. (N.U.J.B.) 

6. Starting from metallic zinc, describe carefully how you would prepare 
reasonably pure specimens of (a) zinc sulphate; (&) zinc carbonate. 

The zinc sulphate crystals are found to contain 43.9% of water of 
crystallisation. Calculate the number of molecules of water in combination 
with one molecule of the anhydrous salt. (O.) 

7. Explain clearly what is meant by a 'normal' salt, an acid salt, and a 
peroxide, giving two examples of each. 

Describe experiments by which you could prove that red lead contains 
lead and oxygen, but no other elements. (L.) 

8 Explain with examples what is meant by the basicity of an acid. 
2 grams of oxalic acid which crystallises with two molecules of water are 
heated with concentrated sulphuric acid. What volume of gas is formed 
at 13°C and 570 mm? If this gas were allowed to stand over potash, what 
volume of gas would remain at s.t.p. (L.) 

9 How would you prepare (a) crystalline sodium carbonate starting 
from sodium hydroxide; (6) lead(II) sulphate starting from metallic lead; 
(c) sodium hydroxide solution starting from sodium carbonate? (N.U.J.B.) 

10 You are supplied with solutions A, B, C, D and E, said to have pH 
values of (A) 2, (B) 7, (C) 7.5, (D) 13, (E) 6.5. Classify these solutions as 
neutral, slightly or strongly acidic, slightly or strongly alkaline. Describe 
a practical method by which you would verify one of the pH figures 
quoted. Which solution would be most likely to liberate hydrogen with 
magnesium powder? Which other solution would be most likely to liberate 
hydrogen with aluminium powder? Write equations for both reactions. 



Chapter 17 



THESE oxides are the ones which correspond to the commonly 
occurring salts. 











Oxides of these metals are 
^soluble in water forming 

Oxides of these metals can 
be made from the metal 
by the action of nitric 
acid and then heat. 
(Al excepted) 

} Oxides of these metals 
decompose when heated. 

Oxides of these metals not 
^reduced to metal by 

Aluminium oxide, AL0 3 

This is a white solid. It is most conveniently prepared by first 
adding dilute ammonia to an aluminium salt solution. This pre- 
cipitates aluminium hydroxide. 

Al,(SO«) 3 + 6NH 4 OH -> 2AJ(OH) 3 + 3(NH«) t S0 4 

The precipitate is then filtered, washed, dried and heated. 

2Al(OH) s -»• Al,0 3 + 3H,0 

If prepared at the lowest temperature possible, it shows both basic 
and acidic properties: 

basic A1 2 3 + 6HC1 -> 2A1C1 8 + 3H,0 

acidic Al,O a + 2NaOH + 3H 2 0-»- 2NaAl(OH)« 



If strongly heated, it passes into a form which is insoluble in both 
acid and alkali. 

Uses. The most important form of this oxide is bauxite, AlgO s . 
2H g O, from which the metal is extracted, p. 477. It also occurs in an 
impure form as emery and is used as an abrasive. 

Coloured by the presence of impurities, this oxide occurs as the 
gems, ruby (iron and titanium), sapphire (chromium) and amethyst 

Summary of preparation of the normal oxides of some common heavy 
metals from the metals or their soluble salts 
Metal (Pb, Cu, Mg, Zn) 

dilute nitric acid 

nitrate of metal 
in solution . 

solid nitrate 



add NaOH solution 1 

< • 

hydroxide of metal 

as precipitate; 

filter; wash; dry 

add Na 2 C0 3 solution 1 


carbonate of metal 

as precipitate; 

filter; wash; dry 





oxide of metal 
(and carbon dioxide) 

oxide of metal oxide of metal 

(+ nitrogen (and water) 

dioxide and 

Oxides of sodium and potassium, Na.X) and K 2 

These oxides are not much used in the laboratory or prepared in 
the usual course of experiment. They react vigorously with water to 
form sodium and potassium hydroxides. 

NaaO + H,0 -*■ 2NaOH or O*- + H s O ->- 20H" 

Calcium oxide (lime, quicklime), CaO 

It is made in industry by the action of strong heat upon limestone, 
calcium carbonate, the latter being placed in a kiln. 
CaC0 3 ->CaO + CO» 

'The hydroxide and carbonate may also be obtained by adding sodium 
hydroxide or sodium carbonate solution to a solution of any soluble salt of the 



Very large quantities of calcium oxide are made in this way (Fig. 
71). In the laboratory, calcium oxide can be made by placing a piece 

of marble in a crucible and 
heating it strongly in a gas- 
heated muffle furnace. A high 
temperature is required and 
an hour or so is necessary to 
complete the action. 

Class preparation of quick- 
lime. Make a loop in a stout 
iron wire large enough to 
hold a piece of marble about 

0t i¥f£/° r the size of a P 63 - Place the 
piece of marble in the wire 

and so arrange it on a tripod 

that when a Bunsen burner 

is placed underneath it, the 

marble is just above the inner 

cone of the roaring Bunsen 

flame. (See Fig. 72.) Leave 

in this position for 5-10 minutes and then allow to cool until the 

solid can be comfortably held in the fingers. 
The original substance (calcium carbonate) and the final product 

(calcium oxide) are very similar in appearance, both being white 

solids. The difference between them can be readily shown. 




Lumps oF 


Fio. 71. 

Continuous process for the 
manufacture of lime. 


Iron Wire 


Fig. 72. 
Calcining marble. 

(a) By the action of water. 
(i) On calcium carbonate— no action. 

(ii) On calcium oxide. Add water a drop at a time to the piece of 
calcium oxide in a dish. Great heat is developed (there may be his- 

sing as the water drops on the calcium oxide), steam is formed and the 
oxide cracks and puffs up and finally crumbles to a powder about 
three times as bulky. This is slaked lime (calcium hydroxide). 

CaO + H g O-»-Ca(OH) t 
slaked lime 

or 0»- + H t O-»-20H- 

AUow more water to fall on to the slaked lime until there is no 
further action. If desired, at this stage, the mixture of slaked lime and 
solution can be filtered and the filtrate shown to be lime-water by 
expelling air from the lungs (containing carbon dioxide) through a 
glass tube into the solution. 
(b) By the action of dilute hydrochloric acid. 
(i) On the calcium carbonate. Effervescence is seen and the marble 
finally disappears. Carbon dioxide is evolved, which, if passed into 
lime-water, turns the latter turbid. 

CaCOg + 2HC1 ->- CaCl* + H,0 + CO, 
or CO g *- + 2H + -> H*0 + CO s 

(ii) On calcium oxide. No evolution of carbon dioxide. The calcium 
oxide will first give a similar action to (a ii p. 244) (slaking) but will 
give finally a colourless solution of calcium chloride. 
CaO + H t O^-Ca(OH) 8 
Ca(OH) a + 2HC1 -> CaCl 2 + 2H s O 
basic acid salt water 


Properties of calcium oxide. Calcium oxide is a white solid. It is 
very refractory; that is, it will not melt even when heated to a very 
high temperature. It merely becomes incandescent and gives out a 
powerful light. It was at one time used for this purpose (lime-light). 
It reacts vigorously with water (see above) to form slaked lime, 
which is an alkali. Its solution in water is called 'lime-water'. Since 
quicklime is basic in character and hygroscopic (that is, it absorbs 
water) it is used to dry ammonia gas. It is used, after slaking, in the 
building trade to make mortar, for manufacture of sodium hydroxide 
(p. 254), and for a very great number of operations needing a cheap 

Zinc oxide, ZnO 

This compound is a white powder (yellow when hot) made in in- 
dustry by distilling zinc and burning the vapour at a jet. 

2Zn + O g ->2ZnO 
or 2Zn + O a ->- 2(Zn* + 2 ") 

It is made in the laboratory from zinc by dissolving the metal in 



dilute nitric acid, evaporating the zinc nitrate solution so formed 
to dryness and heating the residue strongly. 

3Zn f 8HN0 3 ->- 3Zn(N0 3 ) 2 + 4H 2 + 2NO 
2Zn(N0 3 ) 2 -»- 2ZnO + 4N0 2 + 2 

Zinc oxide is amphoteric (see p. 259). 

Zinc oxide (and the oxides already described) cannot be converted 
into the metal by heating the oxides in a stream of hydrogen. 

Zinc oxide, as 'zinc white', is used as a base in the paint industry. 
Although its 'covering power' is not so good as that of white lead 
paint, it does not tarnish, as does a lead paint, on exposure to air 
which contains hydrogen sulphide. The reason is that zinc sulphide 
is white whereas lead sulphide is black. If there is any white paint in 
or near your laboratory, it will probably be a zinc paint. Otherwise 
the hydrogen sulphide fumes would rapidly turn it dark brown or 

lron(III) oxide, Fe 2 3 

(Iron(II) oxide is not important.) This compound is a red powder 
known as 'jewellers' rouge'. It is used for polishing precious stones, 
and as a pigment. 

It is found in the impure state as haematite. Iron(III) oxide is made 
in the laboratory: 

(a) By heating iron(II) sulphate. (Note this action— it is a most 
unusual type.) 

2FeS0 4 -*• F ei O s + SO, + S0 3 

sulphur sulphur 
. . » , dioxide trioxide 

(See p. 419.) 

(b) By heating iron(IIT) hydroxide strongly. 

2Fe(OH) 3 ->-Fe 2 3 + 3H a O 
Iron(lII) oxide is also the product formed if iron(II) hydroxide is 
heated strongly in the air. All iron(II) compounds tend to become 
oxidised to lron(III) compounds by the oxygen of the atmosphere. 

It has the usual properties of an oxide. It can be reduced to metallic 
iron by being heated in a stream of hydrogen or carbon monoxide. 
Fe 2 O a + 3H„ -»• 2Fe + 3H 2 
Fe 2 O s + 3CO -*■ 2Fe + 3CO a 

Tri-iron tetroxide (magnetic oxide of iron), Fe 3 4 

This compound may be prepared by passing steam over red-hot 
iron (p. 265) or by burning iron in oxygen (p. 277). 



It occurs naturally as magnetite and as such is a natural magnet or 
On heating it in a stream of hydrogen it is reduced to iron (see 

d. 158). 

Fe 3 4 + 4H 2 ^ 3Fe + 4H a O 

Lead(U) oxide (litharge), PbO 

This is a yellow powder. It can be made in the laboratory by heating 
lead dioxide, red lead, lead nitrate, lead carbonate or lead hydroxide. 
(For details see compounds concerned.) It is best made from the 
metal by the action of nitric acid with subsequent evaporation, and 
heating of the lead nitrate. When prepared in the laboratory it usually 
ruins the test-tube in which it is prepared by fusing with the glass. It 
is, in fact, used to make lead glass— a glass of very high refractive 

Although lead(II) oxide can be considered a typical base, the only 
common acid in which it will readily dissolve is nitric acid. The reason 
why it does not react quantitatively with the others is a purely 
mechanical one. 

Action of lead(ir) oxide on dilute sulphuric or hydrochloric acids. 
Lead(II) chloride and lead(II) sulphate are not formed quantitatively. 
These two substances are almost insoluble in water and, as the action 
can only proceed on the outside of a particle of oxide, the lead(II) 
chloride or sulphate forms as a layer on the outside. This layer of 
chloride or sulphate is not permeable to the acids and hence action 
stops before any appreciable amount of the salt has been formed. 
(For preparation, see under 'Chlorides' and 'Sulphates', and on 
p. 236.) 

Lead(II) oxide is easily reduced to grey metallic lead by heating it 
in a stream of hydrogen, coal-gas or carbon monoxide. 
PbO + H a -> Pb + H a O 
PbO + CO -* Pb + CO a 

Lead(II) oxide is also an amphoteric oxide dissolving in alkalis to 
form plumbites. 

NaOH + PbO + H a O -> NaPb(OH) 3 or Na + Pb(OH) 3 " 


Copper(ir) oxide (black copper oxide), CuO 

This is made in the laboratory by several methods which are given 

with full experimental details on p. 243. Copper oxide is hygroscopic, 

absorbing moisture from the air. It is a basic oxide and dissolves 

readily in warm dilute mineral acid, forming copper(II) salts, e.g., 

CuO + H 2 SO«-> CuSO« + H t O 



If a solution of copper(II) oxide in concentrated hydrochloric acid 
is gently boiled with clean copper turnings in a fume-cupboard, 
copper(II) chloride is reduced to copper(I) chloride (cuprous chlor- 
ide). This precipitates (white) when the mixture is poured into cold, 
boiled-out water (which prevents re-oxidation). 

CuO + 2HC1 -*• CuCl„ + H 2 

CuCU + Cu -*-2CuCI 
By a similar reduction, a copper(II) sulphate solution, mixed with 
sodium hydroxide solution and sodium potassium tartrate (to pre- 
vent precipitation of copper(II) hydroxide), precipitates red copper(I) 
oxide, or cuprous oxide, when warmed with glucose (a mild reducing 

2Cu*+ + 40H- + C 8 H 12 O a ->- Cu 2 + 2H 2 -f QH 12 7 
Heated copper(II) oxide is reduced to copper by hydrogen or carbon 

CuO + H 2 ->Cu + H a O or CuO + CO -*■ Cu + CO, 
Like manganese(IV) oxide, copper(II) oxide catalyses the decom- 
position of potassium chlorate by heat. 

Mercury(II) oxide, mercuric oxide, HgO 

This red oxide yields, when heated, a mirror of mercury on the 
cooler sides of the test-tube, with oxygen evolved. 

2HgO->2Hg + 1 


The higher oxides arc oxides which contain more oxygen per 
molecule than the corresponding basic oxide. 

Sodium peroxide, Na 2 2 

This is made by heating sodium in excess of oxygen. 

2Na + 2 -»- Na 2 2 

It is a yellow powder and is a vigorous oxidising agent; it should 
never be allowed to come into contact with damp organic matter. 
With water, it liberates oxygen. It is used in confined spaces, e.g., a 
submarine, where men are working, because it absorbs carbon 
dioxide and liberates oxygen at the same time. 

2Na 2 2 + 2H a O -»- 4NaOH + 2 
2Na 2 2 -|- 2C0 2 — »• 2Na 2 C0 3 + 2 

Sodium peroxide is a true peroxide, yielding hydrogen peroxide with 
dilute acids. 



Lead dioxide (LeadfTV) oxide), PbO, 

This oxide is a dark-brown powder and can be made in the labora- 
tory in the following manner: 

Into a beaker put some dilute nitric acid and warm it. By means of 
a spatula, add red lead a little at a time. Care must be taken not to 
add too much red lead or it will contaminate the product. As the red 
lead reacts with the nitric acid, a brown powder is precipitated and 
lead nitrate is formed in solution. The mixture is filtered, the residue 
in the filter-paper is washed two or three times with hot distilled 
water and is allowed to dry on the filter-paper, from which it may 
then be shaken. 

Pb 3 4 + 4HN0 3 -»- Pb0 2 + 2H 2 + 2Pb(NO,), 
red lead lcad(IV) oxide lead nitrate 

The properties of lead dioxide are summed up by the following 
experiments : 

(i) Action of heat. Heat a little lead(IV) oxide in a test-tube and 
hold a glowing splint in the mouth of the test-tube. The splint is re- 
kindled, showing the presence of oxygen. Litharge (lead(II) oxide) 
remains as a yellow solid, often fused into the glass. 
2Pb0 2 -*2PbO + 2 

(ii) Action on concentrated hydrochloric acid. Warm a little lead(IV) 
oxide with concentrated hydrochloric acid. A greenish-yellow gas 
which bleaches litmus is evolved (chlorine) and a white (often dis- 
coloured) solid, lead chloride, may be seen in the test-tube. 
Pb0 2 + 4HC1 -► PbCl. + 2H 2 + Cl 2 

(Compare manganese(lV) oxide.) 

It will be seen from the above that lead(IV) oxide is an oxidising 
agent. If warm, it is converted by sulphur dioxide into lead sulphate 
(a white solid) and the mass glows as combination takes place. 

PbO, + S0 2 -*■ PbS0 4 

(iii) Action of hot concentrated sulphuric acid. Add lead(IV) oxide 
to concentrated sulphuric acid in a test-tube and warm gently. Effer- 
vescence occurs, oxygen is evolved (test as in [i] above) and a white 
precipitate of lead sulphate is left. 

2Pb0 2 + 2H 2 SO«-> 2PbS0 4 + 2H 2 + O, 

Red lead, triiead tetroxide, Pb 3 4 

Red lead is prepared by heating lead(II) oxide for some time in the 
presence of air at a temperature of 450°C. 

6PbO + 2 — »- 2Pb 3 4 



This compound, in many chemical properties, acts as though it 
consists of leadQI) oxide and Iead(IV) oxide. For example, in the 
experiment in preparing Iead(lV) oxide (p. 249): 

Pb 3 Q« -f 4HNQ 8 -*■ 2Pb(NO a ) a + 2H 2 + PbO, 

Compare 2PbO + Pb0 2 + 4HN0 3 
base I acid 

Or, in the action of heat: 
Pb 3 4 

> 2Pb(N0 3 ) 2 + 2H a O + PbO, 
salt water 
unchanged S 

• 3PbO + O 

• 2PbO + PbO + O 

2PbO + Pb0 2 - 

or more correctly : 

2Pb a 4 —*■ 6PbO + O a 

Hence the action of heat on any oxide of lead is to leave lead(II) 

Red lead reacts with concentrated hydrochloric acid or concen- 
trated sulphuric acid when warmed to produce the same observed 
effects as lead(lV) oxide (see above). 

Pb 3 4 + 8HC1 -»■ 3PbCl 2 + 4H 2 + Cl 2 
2Pb 3 4 + 6H 2 S0 4 ->- 6PbS0 4 + 6H 2 + O a 
Red lead has been used for a long time as a pigment (the old name 
for red lead, 'minium', gave the name 'miniature' to that type of 
picture) and is used today, with oil, as a jointing material for gas and 
water pipes and in the manufacture of glass. 

Manganese dioxide, Mn0 2 (Manganese(TV) oxide) 

See Oxygen (p. 273) and Chlorine (p. 355) for the common labora- 
tory uses of this substance. It is also used in the glass industry. 

Questions on this chapter will be found on page 261. 

Chapter 18 


Hydroxides of sodium 
and potassium not de- 
composed by heat. 

Decomposition into 
oxide and water when 
hydroxides of these 
metals are heated. 

K "I Thehydroxidesof these metals are 
Ca > soluble in water and are alkalis. 
Na J (Also ammonium hydroxide.) 

Al "1 These metals form hydroxides'] 
Zn which are insoluble in water. They 
Fe [ are also amphoteric, excepting 
Pb . the two hydroxides of iron. 
Cu Hydroxide is insoluble in water. 
Hg "I Hydroxides of these metals do not 
Ag \ exist. 

Metallic hydroxides are electrovalent compounds, composed of 
metallic ions, which are positively charged, and hydroxyl ions, OH~. 
The number of OH - ions associated with one metallic ion is equal to 
the valency of the metal, e.g., 

Na + OH~; sodium is univalent 
Ca 2f (OH _ ) 2 ; calcium is divalent 
The metallic hydroxides form a very important series of compounds. 
The soluble hydroxides (that is, the alkalis) are particularly im- 
portant. They have many uses both in the laboratory and industry. 

Aluminium hydroxide, Al(OH) 3 

This is a colourless, gelatinous solid. It can be precipitated by 
adding dilute ammonia to a solution of an aluminium salt. The preci- 
pitate is then filtered, washed and allowed to dry. 

Al 2 (S0 4 ) a + 6NH 4 OH->2Al(OH) 3 J + SfNHJjSO, 

Amphoteric nature. If freshly prepared, the hydroxide is soluble in 
dilute mineral acid, forming a salt and water only, and so is basic. 

Al(OH) 3 + 3HC1 -*■ A1C1 3 + 3H a O 



It is also soluble in caustic alkali solution, again forming a salt and 
water only, and so is acidic. 

AI(OH) 3 + NaOH-* NaAl(OH) 4 
Showing both basic and acidic character, it is amphoteric. 

From the above it is obvious that caustic alkali solution, when 
added to an aluminium salt solution, will react in two stages: 

(1) The production of a colourless, gelatinous precipitate of 
aluminium hydroxide. 

Al 2 (SO,) 3 + 6NaOH->2AI(OH) 3 j. + 3N a9 S0 4 

(2) When the alkali is added in excess, this precipitate redissolves 
to give a colourless solution of sodium aluminate. 

AI(OH) 3 + NaOH -»• NaAl(OH) 4 
Uses. This hydroxide has the property of readily absorbing colour- 
ing matter, e.g., if precipitated in a solution containing litmus, it will 
give a blue 'lake' as precipitate. This makes it useful as a mordant for 
certain dyes. It will also carry down bacteria from water and finds a 
use in treating sewage. Precipitated in the fibres of cloth, it makes the 
cloth waterproof. 

Ammonium hydroxide, NH 4 OH 

This is not strictly speaking a metallic hydroxide since the radical 
(NH 4 ) is not a metal, nor has it been isolated. Ammonium hydroxide 
solution is made by dissolving ammonia gas in water (see Ammonia, 
p. 431, for fuller treatment). 

NH 3 + H t O->-NH 4 OH 

ammonIA water ammonlUM hydroxide 
It will turn red litmus blue and precipitate metallic hydroxides from 
solutions of salts of the metal. 

FeCl 3 + 3NH 4 OH Fe(OH) 8 J + 3NH 4 C1 

It forms salts by the neutralisation of acids, for example, 

HNO a + NH 4 OH -> NH 4 NO g + H,0 
ammonium nitrate 

Action of heat. If ammonium hydroxide solution is heated, am- 
monia gas and steam are evolved. 

NH 4 OH->-NH 3 + H,0 
Potassium hydroxide (caustic potash), KOH 

This is made in a similar manner to that used for caustic soda. 
Their properties are almost identical. Note that caustic potash is 
much more soluble in alcohol than is caustic soda, and is used in in- 
dustry to make soft soaps. 



Since the properties are so nearly identical, the cheaper sodium 
hydroxide is almost always used in preference. 

Sodium hydroxide (caustic soda), NaOH 

Sodium hydroxide is produced (at the cathode) by the industrial 
electrolysis of sodium chloride solution (brine). The anode product is 
chlorine. Since these products can react together, 

20H- + Cl 2 -> CI- + OC1- + H 2 
they must be kept apart. In one form of cell (developed by Castner, 
Kellner and Solvay) this separation is effected by the use of a circu- 
lating mercury diaphragm (Fig. 73). The circulation of the mercury 

carbon chlorine spray 

jtj ||| 




hydrogen, water 


J ^M^M 





hydroxide sodium steel 

solution amalgam grids 

Fio. 73. 
Manufacture of sodium hydroxide. 

is shown heavily shaded. In the upper (brine) cell, sodium ion is dis- 
charged and sodium enters the mercury forming sodium amalgam. 
This occurs because hydrogen has a high over-voltage (0.78 volt) at 
a mercury cathode and so does not discharge. Chlorine (an important 
by-product) is liberated at the anode. 

At cathode At anode 

Na+ + e~ -> Na Cl~ - e~ -*• $CU 
Sodium amalgam, flowing into the lower (soda) cell, encounters a 
flow of distilled water in contact with steel grids on which hydrogen 
has only a very low over-voltage. Sodium hydroxide is formed and 
hydrogen is liberated. 

2Na + 2(H*OH-) ->■ 2(Na + OH) + H a 





Much sodium hydroxide is supplied to users (after concentration by 
steam heat under reduced pressure) as a 47% solution. Further 
evaporation by direct fire heat will give molten sodium hydroxide 
which can be cast into solid pellets or sticks. 

The high capital charge on the mercury used adds considerably to 
the cost of this process and rival cells using asbestos separating 
diaphragms, e.g., the Gibbs cell, have been tried. Here, the advantage 
of cheapness is offset by the fact that sodium chloride penetrates the 
diaphragm and pollutes the product. 

Uses of sodium hydroxide 

These uses are very extensive. The following are representative. 

(a) in soap manufacture. A fat or oil is boiled with aqueous sodium 
hydroxide solution (p. 202). 

(b) in the purification of bauxite. This purification yields pure 
alumina, Alj0 3 , for extraction of aluminium (p. 477). 

(c) in the preparation of phenols and cresols from coal-tar. These 
compounds are feebly acidic and are extracted from the tar by sodium 
hydroxide solution, 

e.g., phenol C,H s OH + NaOH ->• C 8 H s ONa + H s O 

(d) in the textile industry. Sodium hydroxide is used in bleaching 
and dyeing processes, in rayon manufacture and in mercerising 
cotton to give it a silky sheen. 

Illustration of Kellner Solva y cell 

Fit up a burette with side tube as shown in Fig. 74. Adjust the taps 
of the two burettes so that mercury flows at approximately the same 
rate from each, and have a sufficient store of mercury in the upper 
burette for about 20 minutes' flow. Connect the electrodes to a 12-volt 
battery (or equivalent supply of electricity), drop a piece of red litmus 
paper into water in the small beaker and arrange a piece of damp 
litmus paper so that gases from the anode come into contact with it. 
After a few minutes, the piece of litmus paper in contact with the 
anode gas is bleached (showing the presence of chlorine) and the 
litmus paper in the beaker has turned blue showing the presence of 
an alkali. To show the evolution of hydrogen from the amalgam, fill 
the narrow tube with mercury and run oft* the bulk of the mercury 
into water. A rapid evolution of hydrogen then takes place. 

Preparation of sodium hydroxide by Gossagc's method 

Half-fill a beaker (of about 600 cm 3 capacity) with water and add to 
it 30 g sodium carbonate as washing soda. Heat the solution almost 
to boiling and adjust the heating to keep it so. Add calcium hydroxide 


Flowing mercury 



Fig. 74. 
Kellner-Solvay cell. 

(slaked lime) a little at a time and stir the mixture. A white solid 
forms on the bottom of the beaker. This is mainly calcium carbonate 
together with some unattacked slaked lime. 

NajCOa + Ca(OH), 

sodium calcium 

carbonate hydroxide 

* CaC0 3 + 2NaOH 

calcium sodium 

carbonate hydroxide 

Filter off a portion of the contents of the beaker and, when the 
filtrate gives no effervescence with dilute hydrochloric acid, it is clear 
that no sodium carbonate remains in it and the double decomposi- 
tion is complete. 



Filter the whole of the mixture and the liquid is a dilute solution of 
sodium hydroxide. The solid can be obtained from it by evaporation 
to dryness, preferably in an iron dish. 

Slaked Lime 

I07o solution of 
Washing soda 
crystals in 
' water 

Fio. 75. 
Preparation of sodium hydroxide from sodium carbonate. 

Preparation of sodium hydroxide by the action of sodium on water 

This is described on p. 196. 

A dilute solution of sodium hydroxide can be obtained by placing 
very small pieces of sodium on water. 

2Na + 2H g O -*■ 2NaOH + H 2 

Properties. Sodium hydroxide is a white solid. It is very deli- 
quescent, has a soapy feel, and will corrode the skin. Care should be 
used in handling the solid and its solution. 

Action on exposure to air. Leave a small piece of sodium hydroxide 
on a watch-glass exposed to the atmosphere for a few days. There is 
formed a solution of sodium hydroxide, showing the solid to be deli- 
quescent. Finally, a white crystalline crust of sodium carbonate is 
formed by the action of the carbon dioxide of the atmosphere. 

Action with water. All deliquescent substances are very soluble in 
water. As sodium hydroxide dissolves, a great amount of heat is 
liberated and the solution should always be made in a thin vessel, 
otherwise the glass may be cracked. 

Action of sodium hydroxide as an alkali (that is, a soluble base). The 



solution is strongly alkaline to litmus, turning it blue. Sodium 

hydroxide will react with acids to form salts (see p. 233), for example 

with hydrochloric acid : 

NaOH + HC1 -*• NaCl + H„0 
sodium dilute sodium water 
hydroxide hydro- chloride 
solution chloric 

or OH- + H+ ^ H 2 

Action of gaseous acidic oxides on alkalis. As a general rule, 
if a gaseous oxide (or hydrogen sulphide) is bubbled into any 
alkaline solution, the final product is the corresponding acid 
salt. (As a rule the normal salt is first formed but it is diificult 
to isolate in most cases.) 

Thus when sulphur dioxide, hydrogen sulphide or carbon 
dioxide are bubbled for some time into potassium hydroxide 
solution, sodium hydroxide solution, lime-water or ammon- 
ium hydroxide, the product is the acid salt; for example: 

CO g + NaOH -*■ NaHCO a 

sodium hydrogen 
H 8 S + KOH ->• KHS + H a O 
2SO* + Ca(OH) 2 -»- Ca(HS0 3 ) 9 
calcium hydrogen 

The normal salt can usually be made by the addition of the same 
weight of alkali as was taken at the commencement of the operation, 
for example: 

H 8 S + KOH -> KHS + H g O 
KHS + KOH -»■ K,S + H a O 

(See p. 400 for experimental details.) 

Sodium hydroxide as a hydroxide-former. The majority of the metal- 
lic hydroxides are insoluble in water, and as we shall see in the 
remainder of this chapter, the usual method of making a hydroxide is 
to add sodium hydroxide solution to the solution of a soluble salt of 
the metal, and isolate the product as described in Chapter 18. 

Action of sodium hydroxide on ammonium salts. If any ammonium 
salt is boiled with sodium hydroxide solution ammonia gas (turns red 
litmus blue) is liberated. 

NH«C1 + NaOH -> NaCl + H,0 + NH, 
or NH 4 + + OH" -* NH 3 + H ,0 



Calcium hydroxide (slaked lime), Ca(OH) 8 

This is a white solid not nearly so soluble in water as sodium or 
potassium hydroxide. 

Preparation. It is prepared commercially and in the laboratory by 
the action of water on calcium oxide (which is obtained by the action 
of heat on limestone). 

CaO + H,0 -»- Ca(OH) 8 
(See p. 244 for experimental details.) 

Properties. Slaked lime is a white solid. Lime-water is formed when 
slaked lime is dissolved in water, but its solubility is only very small 
(about 0.14 g in 100 g water at 20°C). 

This solution is, however, definitely alkaline and will give the usual 
reactions of an alkali. 

Action of heat on calcium hydroxide. Strong heat will convert slaked 
lime into quicklime. In the laboratory the experiment is performed in 
a fireclay furnace by the identical method used to convert calcium 
carbonate into quicklime (see p. 244). 

Ca(OH) 8 ->■ CaO + H 8 

Action of carbon dioxide on lime-water. Lime-water is used as a test 
for this gas owing to the fact that a white precipitate of chalk is 
formed on bubbling carbon dioxide into lime-water. 

Ca(OH) 8 -f CO s -»- CaC0 3 + H,0 

A milkincss is produced which is caused by small particles of solid 
chalk. Further passage of carbon dioxide produces calcium hydrogen 
carbonate, which is soluble in water (see p. 206). 

Uses. Calcium hydroxide is used in the manufacture of mortar and 
bleaching powder and is also used by farmers to counteract acidity 
on the land. 

Mortar is made by mixing calcium hydroxide, sand and water. The 
setting is due to the giving up of surplus moisture and for this reason 
new houses are always damp for some time. With the passage of 
years some of the calcium hydroxide is converted into calcium car- 
bonate by carbon dioxide from the air. 

Ca(OH) 2 + CO s -»- CaCO, + H a O 

Hence, an old mortar will give an effervescence with dilute acids, 
whereas a new mortar will not. 

CaC0 3 + 2HC1 ->- CaCl 2 + H 8 + CO a 
Cement is made by strongly heating limestone and clay together. 
The cement is powdered and, after mixing with a suitable amount of 



water, sets to a mass of interlacing crystals of great strength. Some 

cements will set even under water. When mixed with gravel or brick 

rubble, cement forms concrete, and is used with water in the same 


Zinc hydroxide, Zn(OH) 8 

This is a white powder. It is formed as a white precipitate when 
sodium hydroxide solution is carefully added to the solution of a 
soluble zinc salt (usually zinc sulphate). 

ZnS0 4 + 2NaOH ->- Zn(OH) 8 -1- Na 8 SO, 
or Zn 2+ + 20H- -»- Zn(OH), 

The white precipitate, if required, may be filtered off, washed with 
hot distilled water and dried on a porous plate. 

Amphoteric nature of zinc hydroxide. Care must be taken in the 
above experiment not to add excess sodium hydroxide solution, since 
zinc hydroxide is soluble in excess of caustic alkali solution, forming 
a zincate. 

Zn(OH) 2 + 2NaOH -► Na^nfC-H), 
or Zn(OH) 8 + 2C-H" -► Zn(OH)««- 

Here zinc hydroxide is acting with acidic properties because, with the 
alkali sodium hydroxide, it forms the salt, sodium zincate, and water. 

Like most metallic hydroxides, zinc hydroxide has basic properties 
also, and it reacts, for example, with dilute sulphuric acid to form the 
salt, zinc sulphate, and water. 

Zn(OH) 2 + H 8 S0 4 -► ZnSC- 4 + 2H a O 
base acid salt water 

A metallic hydroxide which exhibits both basic and acidic properties 
is said to be amphoteric. The corresponding oxides are also am- 
photeric. . 

Action of heat. Zinc hydroxide is readily converted into zinc oxide 

by the action of heat. 

Zn(OH) 8 -vZnO + H 2 



(yellow hot, 

white cold) 

Hydroxides of iron, Fe(OH)i!, Fe(OH) 8 

Iron forms two hydroxides, iron(II) hydroxide and iron(III) 

hydroxide. . , 

Iron(U) hydroxide is precipitated by the action of sodium hy- 
droxide solution on iron(II) sulphate solution. 

FeS0 4 + 2NaOH->- Fe(OH) 8 | -f Na s S0 4 
or Fe 2+ + 20H--»-Fe(OH) 8 



If air is excluded the substance is white, but under ordinary con- 
ditions it is seen as a dirty-green gelatinous precipitate. It is trouble- 
some to isolate the solid, as it is so readily oxidised by the air. Make 
the precipitate in the way indicated above and allow it to stand or 
warm it. It will be seen to turn reddish brown because it is oxidised 
by the oxygen of the atmosphere to reddish-brown iron(III) hy- 

4Fe(OH) 1 + O s + 2H a O -*■ 4Fe(OH) 3 

For the same reason, on heating ironfjl) hydroxide strongly, iron(Ili) 
oxide remains. 

Jron(IJI) hydroxide is precipitated as a reddish-brown gelatinous 
precipitate by adding an alkaline solution (for example, sodium 
hydroxide solution) to a solution of iron(III) chloride. 

FeCIg + 3NaOH -> Fe(OH) 3 + 3NaCI 
or Fe 8 + + 30H- —*■ Fe(OH) 3 

To purify it, filter off, wash well with hot distilled water and allow it 
to dry. 

On heating the hydroxide strongly, iron(III) oxide is obtained as 
a red powder. 

2Fe(OH) s -»- Fe s 3 + 3H s O 

Ironfjl) and iron(III) hydroxide dissolve in dilute acids with the 
formation of salts and water, thus indicating their basic nature. 

Lead(D) hydroxide, Pb(OH), 

This is formed as a white precipitate by adding sodium hydroxide 
solution cautiously to lead nitrate solution. 

Pb(N0 3 ) 2 + 2NaOH -> Pb(OH) 2 + 2NaN0 3 
or Pb a+ (aq) + 20H~(aq) -► Pb(OH) 2 (c) 

It can be obtained as a white powder by filtering off, washing with hot 
distilled water, and allowing it to dry. 

Like zinc hydroxide it is amphoteric and will dissolve in acids to 
form lead(II) salts and in alkalis to form plumbites. 

Pb(OH) 2 + 2HNO s -► PbfNO^j + 2H 2 
Pb(OH) 2 (c) + OH-(aq)-> Pb(OH) 3 -(aq) 

plumbite ion 

On strongly heating lead(II) hydroxide, Iead(II) oxide remains as 
a yellow powder. 

Pb(OH) 2 -»-PbO + H 2 

Copper(II) hydroxide, Cu(OH) 2 

This is formed as a blue gelatinous precipitate on adding sodium 
hydroxide solution to cold copper sulphate solution. The substance 



can be obtained as a blue powder by filtering off the precipitate, 
washing it several times with cold distilled water and allowing it to 


CuSO« + 2NaOH ->- Cu(OH) 2 + Na 2 SO, 
or Cu*+(aq) + 20H "(aq) -»• Cu(OH) 2 (c) 

It is soluble in dilute acids forming solutions of copper(II) salts. For 

Cu(OH) 2 + H 2 S0 4 -> CuS0 4 + 2H g O 
On warming the hydroxide, it readily loses water and forms black 
copper oxide (see p. 26). 

Cu(OH) 3 -»- CuO + H 2 
Copper(II) hydroxide dissolves in ammonia to form a deep blue 
solution. This blue solution will dissolve cellulose (for example, a 
filter-paper) and this solution of cellulose has been used to make 
artificial silk by forcing the solution through tiny holes into an acid 
solution. The fibres so formed are spun into threads. 


1. Describe a method for Ihe preparation of sodium hydroxide from 
sodium chloride. 

Give an account of the reaction of sodium hydroxide with (a) carbon 
dioxide; (b) zinc; (c) a solution of zinc sulphate. 
Give the equations in each case. (C) 

2. Sulphur dioxide is called an acidic oxide and copperfjl) oxide a basic 
oxide. What is meant by these terms? Give two other examples of each of 
these classes of oxides and describe how you would test an oxide in order 
to assign it to one of these classes. (O. and C.) 

3. What chemical properties distinguish metals as a class? Describe in 
detail three methods by which a metal may be converted inlo its oxide. In 
each case name the metal which you would use. (L.) 

4. How would you obtain from calcium carbonate pure specimens of 
(a) calcium oxide; (b) calcium hydroxide? Describe the properties of an 
aqueous solution of calcium hydroxide, explaining what changes occur 
when carbon dioxide is passed through the solution. (O. and C.) 

5. Name five gases which can be absorbed by sodium hydroxide. Explain 
the reactions which take place on absorption, and in each case give one 
test by which you could recognise the product. (L.) 

6. How may lead(II) oxide and lead(lV) oxide be obtained, starting from 
metallic lead ? 

Describe how you would show by experiment that their composition is 
in agreement with the Law of Multiple Proportions. (L.) 

7. What is the chemical nature of galena and white lead? Starting with 
red lead (PbjOi), how would you prepare a specimen of each of the 
following; (a) metallic lead; (6) lead(II) oxide; (c) lead(TV) oxide? 



8. Explain the meaning of the terms base, basic oxide, hydroxide. 
Classify the following substances: lime, ammonia, caustic soda. Give 
reasons for your classification. (O. and C.) 

9. Define base, hydroxide, alkali, giving two examples of each with 
formula. Indicate briefly how you would prepare (a) sodium hydroxide, 
using sodium carbonate; (b) iron(TIl) hydroxide, using iron(IU) chloride. 

10. What do you understand by (a) a base; (6) an alkali? 

Draw a labelled diagram to show how sodium hydroxide is prepared 
electrolytically. (The diagram should show the nature of the electrolyte 
and all products, the names, signs and materials of the electrodes, and any 
other details essential to the process.) 

Stateconcisely what you would observe when a stick of sodium hydroxide 
is left exposed to the air for a long time. (Equations are not required in this 
question.) (N.U.J.B.) 

1 1 . What are the four chief classes of oxides ? Give one example of each 

How do oxides of metals differ chemically from those of non-metals? 

State briefly how you would prepare from suitable oxides (a) a solution 
of copper(II) sulphate; (6) a dilute solution of hydrogen peroxide. 

What happens when manganese(IV) oxide is added to a solution of 
hydrogen peroxide? (N.U.J.B.) 

Chapter 19 


Dilute sulphuric and hydrochloric acids attack these metals 
\ with the liberation of hydrogen. 
(Al with dilute HC1 only) 













Hydrogen was first recognised by Cavendish (1766). It was called 
'inflammable air', and the name hydrogen {i.e., water producer) was 
given to it by Lavoisier. 


Uncombined hydrogen does not occur in nature to any appreciable 
extent, but the element occurs in vast quantities in a combined state 
in such compounds as water, acids, and many organic substances. 

Industrial preparation of hydrogen 

Hydrogen has acquired much greater importance in recent years 
because of new uses given below. There are now two chief methods 
of manufacture— from hydrocarbons and by electrolysis. 

(1) From hydrocarbons. In recent years, hydrogen has been made 
almost entirely from hydrocarbons. To use the simplest example, 
methane (natural gas) can be passed with steam over nickel catalyst 
at 800°C and 30 atm. 

CH 4 +H t O->-CO + 3He 
The product is mixed with more steam and passed over iron(IH) oxide 




(catalyst) at 450°C. Carbon monoxide is converted to the dioxide with 
further yield of hydrogen. 

CO + 3H 2 + H 2 ->- C0 2 + 4H 2 

Carbon dioxide is dissolved out by water under 30 atm pressure. 

H,0 + CO, ^ H 2 C0 3 

Any remaining traces of carbon monoxide are absorbed under pres- 
sure, by copper(I) formate in ammonia. 

(2) By electrolysis. Hydrogen is obtained as a by-product in the 
electrolytic manufacture of sodium hydroxide from common salt 
(p. 253). 

Where electrical power is cheap, hydrogen can be made by elec- 
trolysis of water containing sulphuric acid. The laboratory version of 
this method is fully discussed on p. 146. 

Uses of hydrogen 

(1) For filling balloons. It is the lightest gas known but has the 
great disadvantage of inflammability. 

(2) In the 'hardening" of oils to make margarine. Oils, e.g., olive oil 
or whale-oil, are heated to 180°C and finely divided nickel is added 
as catalyst. They are then treated with hydrogen at about 5 atm 
pressure. The oil combines with hydrogen and is converted to a fat, 
which is solid at ordinary temperature and is used in the manufacture 
of margarine. In this way, a liquid oil, unacceptable in our diet, is 
'hardened' to an acceptable solid fat and used as a butter-substitute. 

(3) In the conversion of coal to synthetic 'petrol'. A paste of coal- 
dust and oil, containing iron oxide and alkali as catalyst, is sprayed 
into hydrogen at 200 atm pressure and 450°C. The products are 
hydrocarbon gases and a liquid oil. On distillation, this oil yields a 
fraction suitable for petrol. (This process has been made obsolete 
by large, cheap supplies of natural petroleum.) 

(4) In the manufacture of ammonia. This is Haber's Process, des- 
cribed on p. 431. 

(5) In the synthesis of hydrochloric acid (H 2 + Cl 2 — *■ 2HCI) and 
of organic chemicals, e.g., methanol, CH 3 OH. Also in the oxy- 
hydrogen flame for cutting and welding steel. 

Preparation of hydrogen by the action of dilute acids on metals 

Note. Hydrogen explodes violently with air if a spark or 
flame reaches the mixture. For safety, always collect a 
sample of hydrogen and test as described on p. 47 before 
lighting a jet or collecting the gas in bulk. 

Into a flat-bottomed flask or bottle, put some pieces of zinc and 
add dilute sulphuric acid by means of a thistle funnel. There is 



Dilute Sulphuric 
1 Acid 



Fig. 76. 
Preparation of hydrogen. 

effervescence, and a gas is given off which is collected over water, as 
shown in Fig. 76. Zinc sulphate, which is formed, dissolves in the 
water to form zinc sulphate solution. 

Zn + H 2 SO,-»- ZnSO, + H a 
or Zn + 2H 8 0+ -> Zn* + + 2H.O + H, 

Hydrogen from water. Action of steam on heated iron 

This method is very suitable for obtaining hydrogen in quantity. 
The apparatus consists of a long iron tube loosely filled with iron 

Safety tube 

Iron nails and Filings 



Fio. 77. 
Preparation of hydrogen in quantity from steam. 



filings and nails which arc heated to redness by a furnace (Fig. 77). 
Steam is generated by boiling water in a can, and the steam passes 
over the heated iron, forming hydrogen, which is collected over 
water. Any excess steam condenses as it comes into contact with the 
cold water. The iron is converted into tri-iron tetroxide (black oxide 
of iron). 

3Fe + 4H a O -*■ Fe 3 4 + 4H 2 
iron steam tri-iron hydrogen 


This reaction is reversible (see p. 158). 

Hydrogen from alkalis 

Warm sodium (or potassium) hydroxide solution will react with 
zinc, aluminium or silicon to liberate hydrogen and leave a solution 
of sodium (or potassium) zincate, aluminate or silicate. 

Zn + 2NaOH + 2H a O -»- Na a Zn(OH)« + H, 

2A1 + 2NaOH + 6H a O -* 2NaAl(OH) 4 + 3H 2 

Si + 2NaOH + H a O -> Na^SiOa + 2H a 

These methods are not usually used in the 

Test for hydrogen. A mixture of hydrogen and air 
explodes when a flame is applied. 

Properties of hydrogen 

Hydrogen is an invisible gas, neutral to litmus, 
and, if pure, possesses no smell. 

Hydrogen is lighter than air. Hold a gas-jar full 
of hydrogen (B) under a gas-jar full of air (A) and 
take off the cover (see Fig. 78). 
Fio. 78. Count ten. Now apply a lighted splint to the 

gas-jar which previously contained air, and the gas 
in it will burn or explode, showing hydrogen has passed upwards 
into it. On applying a flame to the gas-jar which originally con- 
tained hydrogen, the flame produces no change, showing the gas to 
be air. 

Hydrogen burns in air. Invert a gas-jar of hydrogen and apply a 
lighted splint. The gas burns quietly round the edges of the gas-jar. 
A splint pushed up into the gas will be seen to be extinguished. Hence 
hydrogen will burn in air but will not allow a splint to burn in it. 
The product of the combustion of hydrogen is water. 

2H a +O a -*2H 2 
Hydrogen explodes if mixed with air and a flame is applied. (The 
usual test for hydrogen.) When collecting the hydrogen over water, 



bubble the gas into a gas-jar containing about two thirds of its volume 
of air, passing in hydrogen until the water is all displaced. Apply a 
flame to the mixture and there is an explosion. 
2H a + 1 -»-2H a O 
Hydrogen is a reducing agent. Hydrogen readily reduces the oxides 
of copper, lead or iron to the metal when they are heated in a stream 
of the gas, i.e. (see p. 48), 

CuO + H 2 -*Cu + H a O 
PbO + H a ->Pb + H a O 
Fe 2 3 + 3H a -»- 2Fe + 3H a O 
In these cases, the system of oxide ion and hydrogen makes electrons 
available and so acts as a reducing agent. 

2 --r-H a ->H a O + 2e- 

The metallic ion acts as an oxidising agent by accepting the electrons 
and is reduced to the corresponding metallic atom. 
Cu* + + 2e~ -* Cu (Pb similar) 
2Fe 3+ + 6e- -»■ 2Fe 
Nascent hydrogen. A system of acid and metal which is generating 
hydrogen, e.g., zinc and dilute hydrochloric acid, has a very vigorous 
reducing action. For example, it will reduce iron(Ui) ion in solution 
to iron(H) ion, though gaseous hydrogen passed through the iron(III) 
ion solution does not reduce the ion at all. The reducing action used 
to be ascribed to so-called nascent hydrogen, i.e., hydrogen at the 
moment of its chemical production, which was assumed to be very 
reactive. This conception has been abandoned. The prevailing idea 
now is that, in the presence of the acid, a fairly electropositive metal, 
e.g., zinc, makes electrons available; that is, it acts as a reducing 
agent. The electrons are accepted by an oxidising agent, e.g., iron(UI) 
ion, and it is reduced— in this case to ironfjl) ion. 
Zn -► Zn* + + 2e~ 
2Fe a+ + 2e--*2Fe il+ 
Adding these, Zn + 2Fe s+ -> Zn* + + 2Fe* + 

Sources of the so-called nascent hydrogen in alkaline conditions are 
sodium amalgam and water, or aluminium and sodium hydroxide 
solution. In these cases, also, the mixtures act as reducing agents by 
making electrons available. 

NaHg->-Hg-r-Na + -f-e- 
Al + 40H- -► Al(OH)r + 3e~ 
Action of hydrogen on the halogens. Hydrogen combines with 
chlorine to form hydrogen chloride, 

H a +Cl a -»-2HCl 



with bromine, less readily, to form hydrogen bromide 

H, + Br a ->-2HBr 
with iodine, less readily still. (See Halogens.) 

Isotopes of hydrogen 

This subject is, in the main, beyond the scope of the present book. 
It may be briefly mentioned, however, that hydrogen has at least 
three isotopes— ordinary hydrogen or protium, H, heavy hydrogen or 
deuterium, D, and tritium, T. All these have one electron per atom. 
In protium, the nucleus of the atom consists of a single proton; in 
deuterium, the nucleus contains a proton and a neutron; in tritium, it 
contains a proton and two neutrons. The approximate atomic weights 
are, therefore, H = 1, D = 2 and T = 3. Protium and deuterium 
resemble each other closely in chemical behaviour, deuterium being 
somewhat less reactive. Deuterium oxide is known as heavy water, 
D 2 0. It resembles ordinary water but has a higher density (about 
1.10 gem -3 at room temperature). Tritium is radioactive; it has a 
half-life of about twelve years— that is, in that time, half of the 
tritium is lost by conversion into the products of the radioactive 


Invert a gas-jar of hydrogen over a gas-jar of air (i.e., the hydrogen 
is the uppermost of the two). Remove the plates and leave for ten 

minutes. Plunge a lighted taper 
into each gas-jar separately and 
in both cases there is an explo- 
sion, showing hydrogen to be 
present in quantity in both of 
the gas-jars. Hence some of the 
hydrogen must have descended 
from the upper gas-jar to the 
lower. This seems contrary to 
the result of the experiment 
above which shows the gas to 
be less dense than air. 

This process of filling the 
whole of a vessel in which a 
gas is placed is termed dif- 
fusion. Hydrogen diffuses (i.e., 
spreads throughout the vessel 
Fig. 79. in which it is placed) more 

Coal-gas diffuses more rapidly than air. rapidly than any other gas. 

\ Level 




This can be shown by the simple apparatus depicted in Fig. 

To show the rapidity with which hydrogen diffuses. A porous pot has a 
rubber stopper which fits tightly in its mouth and is connected to a bent 
glass tube containing coloured water (Fig. 79). 

A beaker surrounds the pot and coal-gas is passed into the beaker by 
placing an unlit Bunsen burner under it and turning on the gas tap. On 
doing so, the level of water in the limb directly below the pot begins to fall 
because the coal-gas (which contains about half its volume of hydrogen) 
diffuses through the tiny pores of the pot more quickly than the air diffuses 
out of the pot. This causes an excess pressure of gas in the pot, making the 
level fall. When the level returns to normal, as it will do in a few moments, 
the beaker containing the coal-gas is removed. Now the reverse happens 
and the level of the water in the shorter tube rises. This is because the coal 
gas (now inside the pot) diffuses out more rapidly than the air can diffuse in, 
and causes a lowering of the pressure inside the pot, and hence the level 


Graham found that the rate of diffusion of a gas was inversely propor- 
tional to the square root of its density in constant conditions. For example, 
since the density of oxygen relative to hydrogen (T. and P. constant) is 16, 
Rate of diffusion of hydrogen = V16 = 4 
Rate of diffusion of oxygen VI 1 

That is, hydrogen diffuses four times as rapidly as oxygen. This may be 
applied as a method of rinding the molecular weight of a gas by finding its 
vapour density from its rate of diffusion. 

All gases, diffuse, i.e., they attempt to distribute themselves uniformly 
throughout the containing vessel. For example, if a gas-jar of (brown) 
bromine vapour (density relative to hydrogen = 80) is placed below a gas- 
jar of hydrogen, the bromine diffuses upwards and the hydrogen downwards 
until a uniform brown colour is seen throughout the gas-jar. Similarly (but 
very much more slowly), if a large crystal of (blue) copper sulphate is put 
into a gas-jar which is then about half-filled with water and left to stand 
(covered), the salt will dissolve and its ions will diffuse upwards while water 
molecules diffuse downwards till a uniformly blue solution is obtained. 

The burning of hydrogen in air to form water 

This is the synthesis of water, i.e., the building-up of water from its 

Fit up the apparatus as shown in Fig. 80. 

Hydrogen is generated by the action of dilute hydrochloric acid on 


Zn + 2H+ -> Zn 8+ + H s 

The gas then passes through a U-tube containing calcium chloride in 
order to dry the gas, the hydrogen is burnt at a jet, and the vapours 
are cooled by coming into contact with a can kept cool by water. 

When the apparatus has been set up (use rubber stoppers or very 
well-fitting bark corks), place a test-tube over the jet and collect a 
test-tube full of hydrogen by displacement of air. When this test- 



tube full of gas burns quietly on the application of a flame to it, 
light the jet and allow the flame to burn so that it just does not touch 
the cooled can. 

Bright metal can 


Calcium Chloride 
to dry the. 

■ Zinc and dilute 
Hydrochloric acid 

Fio. 80. 
Synthesis of water. 

Moisture will condense on the can and will drop off into a dish 
which is placed below to receive the liquid. 
The liquid can be shown to be water by the tests described below: 

Slit fortscope 
of vapour \ 

- 100'C If A.P. 
is 760 mm. 

Cotton-wool to prattel 
tht bulb from 

-l/au/d for lest 


1. Action on anhydrous copper sulphate. 

Allow a drop of the liquid to fall on to anhydrous copper sulphate. 
A blue patch on the white solid (with hissing and development of 
heat) proves water is present, but does NOT prove the liquid to be 
pure water. 


CuS0 4 + 5H 2 = CuS0 4 .5H 2 
blue vitriol 


2. Boiling-point (see Fig. 81). 

If the atmospheric pressure is 760 mm, the thermometer should 
register 100°C. 

These two tests together prove the liquid to be pure water. An 
additional, but less convenient test, is to find the freezing-point, 
which should be 0°C at 760 mm pressure. 

2H, + 2 = 2H s O 



1 . What weight of sulphuric acid would be necessary to react with zinc 
to provide sufficient hydrogen to fill 6 gas-jars, each containing 400 cm 3 
at 15°Cand750mm? 

2. Hydrogen is a reducing agent. Illustrate this statement by reference 
to the action of the gas with (a) oxygen; (b) chlorine; (c) copperfM) oxide. 

3. If a mixture of hydrogen and oxygen is passed slowly through a long 
porous tube the gas issuing at the other end does not explode when a flame 
is applied. Explain this. If the same experiment was performed with a 
mixture of carbon monoxide and air, would you expect the gas issuing at 
the end to explode or burn when a flame is applied ? Give reasons for your 

4. Describe two methods by which you could prepare reasonably pure 
hydrogen. . . 

Calculate (a) the volume; (Z>) the percentage composition of the gas 
which remains when equal volumes of air and hydrogen are exploded and 
the products are allowed to cool down to room temperature again. (Air 
contains 21% of oxygen by volume.) (O.) 

5. By what properties could you distinguish between the compounds 
which hydrogen forms with (a) chlorine; (b) sulphur; (c) nitrogen? Sketch 
the apparatus you would employ to prepare and collect a sample of one of 
these compounds. (N.U.J.B.) 

6. How would you prepare dry hydrogen, and show that it forms a 
liquid when burned in air? Draw the apparatus you would use. By what 
experiments could you show that the liquid (a) contained water; (6) con- 
tained nothing but water? Give full details. (N.U.J.B.) 



Chapter 20 












When heated in air these metals oxidise with a readiness 
indicated by the order shown; that is, potassium most 
easily, copper least readily. 

| These metals do not oxidise easily; their oxides yield 
f oxygen on heating. 

Oxygen was first discovered by Scheele in 1772, but Priestley dis- 
covered it independently two years later by heating oxide of mercury. 
He called it 'dephlogisticated air', but Lavoisier called it oxygen 
(acid-producer) because he obtained acids by heating several non- 
metals in oxygen and dissolving the oxides in water. 


Uncombined oxygen exists in the air, forming 21% by volume (or 
23% by weight). The earth's crust consists of almost half of its 

Chlorate 4 

Dioxide I 

Fig. 82. 

Preparation of oxygen from potassium chlorate. 


weight of oxygen in a combined state in the form of water, silicates, 
many metallic and non-metallic oxides, and in the form of 

Laboratory preparation of oxygen from potassium chlorate 

Crush some potassium chlorate crystals (20 g) in a mortar and 
grind with them about one-quarter of their weight (5 g) of man- 
ganese dioxide (manganese(IV) oxide). Place the mixture in a hard- 
glass tube and fit up the apparatus as shown (Fig. 82). Heat the 
mixture and a gas will readily be given off which can be collected 
over water. 

Since oxygen has about the same density as air, it cannot 
be collected by displacement of air. If required dry it can be 
dried by calcium chloride and collected over mercury. 

Potassium chloride is left in the tube. 

2KC10 3 ->-2KCl-}-30 2 
or 2C10 3 -->-2Cl- + 30 2 

Test for oxygen 

It rekindles a glowing splint of wood. This distinguishes 
it from all gases except dinitrogen oxide, N a O. It is distin- 
guished from this gas: 

(a) by having no smell (dinitrogen oxide has a sweet, sickly 

(b) with nitrogen monoxide, oxygen produces brown fumes of 
nitrogen dioxide. 

2NO + 2 ->2N0 2 

Dinitrogen oxide has no effect on nitrogen monoxide. 

Manganese dioxide (manganese(TV) oxide) as a catalyst 

If potassium chlorate is heated alone, it gives off oxygen, but only 
at a fairly high temperature (400°C). If mixed with manganese 
dioxide, the potassium chlorate gives off oxygen at a much lower 
temperature and much more steadily. On analysis of the residual 
mixture, it is found that the amount of manganese dioxide is exactly 
the same at the end of the experiment as it was at the beginning. 

A substance which can alter the rate of a chemical reaction in 
this way is called a catalyst. 

Definition. A catalyst is a substance which, although present in small 
proportions, alters the rate of a chemical reaction, but remains chemi- 
cally unchanged at the end of the reaction. 



Class experiment to show manganese dioxide is a catalyst 

Mix a little manganese dioxide with about four times its bulk of 
potassium chlorate and place in an ignition tube. Into each of two 
other tubes put an approximately equal bulk of manganese dioxide 
and potassium chlorate respectively. Surround each with sand on a 
sand tray so that they are close together and vertical. (See Fig. 83.) 
Commence to heat and test at intervals for oxygen by lowering a 
glowing splint into each test-tube. After about one minute oxygen is 
freely evolved from the mixture, with no signs of gas from either of 
the other tubes. 

In order to be quite certain that the oxygen was coming from the 
chlorate and not merely from the manganese dioxide, it would be 
necessary to show that the weight of manganese dioxide was the same 
after the experiment as before it. This could be done by dissolving 
the chlorate and chloride of potassium in water and recovering the 
manganese dioxide by filtering, washing with water, and drying in a 
steam oven. It would be found that there was no appreciable loss of 
weight. For further discussion of catalysis, see p. 155. 




Dioxide *-s 


Mixture of 

Fio. 83. 
To show manganese dioxide is a catalyst. 

Laboratory preparation of oxygen from hydrogen peroxide 

This is a convenient preparation because it requires no heat. 
Hydrogen peroxide solution (20 vol.) is added, drop by drop, to 
manganese dioxide, which catalyses decomposition of the peroxide. 
Oxygen is collected over water (Fig. 84). 

2H,0,-»-2H 2 0-f O s 

An alternative preparation (in the same apparatus) is the drop by 
drop addition of hydrogen peroxide solution (20 vol.) to potassium 
permanganate in the presence of excess of dilute sulphuric acid. 
Oxygen is liberated until all permanganate is decomposed, by which 
lime the mixture is colourless. 
5H,0, + 2KMn0 4 + 3H,SO«-»- K 2 S0 4 + 2MnSO t + 8H.O + 50, 

20 vol. 









Fig. 84. 
Preparation of oxygen from hydrogen peroxide. 

Properties of oxygen 

Oxygen is a colourless, odourless, neutral gas, is almost insoluble 
in water and has approximately the same density as that of air. It is 
an exceptionally active element, combining vigorously with many 
metals and non-metals, forming basic and acidic oxides respectively: 

• METALLIC OXIDES most of which 

are basic 
in character 

which are 


in character 

Action of oxygen with non-metals to form acidic oxides 

Phosphorus. Place a small piece of yellow phosphorus in a defla- 
grating spoon, warm it until it begins to burn, and then plunge it into 
a gas-jar of oxygen (Fig. 85) into which you have previously poured 
a little blue litmus solution. The phosphorus burns with a dazzling 
flame, emitting white clouds of oxides of phosphorus which dissolve 
in the water to form acids of phosphorus, which turn the litmus red. 

Pi + SOj-^-P^xo 





P 4 + 30„->-P 4 0„ 


On solution in water: 
6H 2 O + P«O 10 -h 

6H,0 + P 4 0, 

-* 4H 3 P0 4 

> 4H 3 PO a 

Fig. 85. 

Phosphorus burning in 

Sulphur. In a similar manner, lower a 
piece of burning sulphur into a gas-jar 
of oxygen containing blue litmus solu- 
tion. Misty fumes of sulphur dioxide 
are given off as the sulphur burns more 
brightly with its characteristic blue flame, 
and this gas dissolves in the water to 
form sulphurous acid, which turns the 
litmus red. 

S + O g — »• SO* 

sulphur oxygen sulphur 

SOs + HjO-i-HjSOa 



Carbon. Perform the same experiment with wood charcoal (car- 
bon). The charcoal burns and emits a shower of sparks, combining 
vigorously with the oxygen to form a colourless gas, carbon dioxide 
which dissolves in the water to form carbonic acid. This is only a 
very weak acid and the litmus is turned claret-coloured but not 
definitely red. 

CM-O t -*-COj 

carbon dioxide 


carbonic acM 

If the above experiment is performed with lime-water in the place of 
litmus solution, the lime-water will become turbid because of the 
formation of a precipitate of chalk. 

Ca(OH),-f CO, 




■* CaC0 3 + H,0 

The above oxides are examples of anhydrides (that is, 
oxides of non-metals which react with water to form acids), 


and it was because of these that Lavoisier gave the gas the 
name oxygen (acid-producer). 

Action of oxygen with metals to form bases 

Magnesium. Lower a piece of burning magnesium ribbon by means 
of tongs into a gas-jar of oxygen. It burns with a more dazzling 
flame and forms a white ash, magnesium oxide. 

2Mg + 2 -»-2MgO 
or 2Mg + 0,->2(Mg 2+ 0*-) 

Here, and in similar cases, oxygen is acting as an oxidising agent by 
accepting electrons (from the metal). 

Iron. Attach a piece of iron wire to the end of a deflagrating spoon 
and dip the end of the wire in sulphur (to start the action). Warm the 
wire in the Bunsen flame until the sulphur begins to burn and then 
plunge it quickly into a gas-jar of oxygen which contains a little 
water. The iron wire burns, giving off a shower of sparks, and finally 
a molten bead of oxide drops into the water. 

3Fe +-20 2 -*-Fe 3 4 

Industrial preparation 

Since oxygen exists to such a large extent in air, it is natural for 
attempts to be made to obtain it from this source. It is not easy to do 
this, since nitrogen is an inert element and cannot readily combine 
with anything and thus leave the oxygen pure. By far the best pro- 
cess for obtaining oxygen industrially is from liquid air. 

Liquid air. Air is first compressed to about 200 atmospheres pres- 
sure, cooled and allowed to escape from a small jet. Expansion cools 
the air further because heat energy is used in separating the molecules. 
The cooled air is allowed to flow away by passing round tubes con- 
taining the incoming compressed air. This cools the incoming air 
and these successive coolings are finally sufficient to liquefy the 
air. On evaporation of the liquid, nitrogen (b.p. 77 K) is first evolved, 
leaving a liquid very rich in oxygen (b.p. 90 K at 760 mm). This is 
a fractional distillation of the liquid air. Oxygen is sold for com- 
mercial use, not as liquid, but as gas compressed in strong steel 
cylinders at about 100 atmospheres pressure. 


(1) As an aid to breathing where the natural supply of oxygen is 
insufficient, for example, in high-altitude flying or climbing, and also 
when anesthetics are administered to a patient. 



(2) In the oxyacetylenc flame, which can be used for welding and 
for cutting even very thick steel plate. The temperature of the flame 
reaches about 2200°C. 

(3) In the new L-D process for making steel (p. 486) A great and 
increasing oxygen tonnage is now used in this way. 

The oxygen molecule 

It is known directly by experiment (p. 115) that 2 volumes of hydrogen 
combine with 1 volume of oxygen to form 2 volumes of steam (in constant 
conditions which do not condense the steam). Applying Avogadro's Law 
to the oxygen and steam, 1 molecule of oxygen yields 2 molecules of 
steam. Since the minimum quantity of oxygen possible in 1 molecule of 
steam is 1 atom (atoms being indivisible for chemical purposes), 1 
molecule of oxygen must supply at least 2 atoms of oxygen and is, 
therefore, at least a diatomic molecule. As additional evidence, no case 
is known in which 1 volume of oxygen produces more than 2 volumes 
of a gaseous compound in constant conditions. This points to an actually 
diatomic molecule, O t , for oxygen. 

Further, the ratio of the two specific heats of gaseous oxygen is (by 
experiment) 1.40, specific heat at constant pressure being greater than 
specific heat at constant volume. This is the recognised ratio for a diatomic 
gas and confirms the diatomicity of the oxygen molecule. 

Classification of oxides 

Four important types of oxide are the following. 

A basic oxide is a metallic oxide which reacts with an acid to 
produce a salt and water only; if soluble in water, it forms an alkaline 
solution, e.g., 

CaO + 2HC1 -»• CaCl a + H s O; CaO 4- H,0 -> (Ca(OH), 

Other examples are: Na 8 0, K 2 (alkalis NaOH, KOH); CuO, 

An acidic oxide is a non-metallic oxide which, when combined 
with the elements of water, produces an acid, e.g., 

SO, + H.O ->- H 2 SO«; P<0,o + 2H 2 -> 4HPO s 
Other examples are C0 2 , SO,, SiO g (acids, H 2 CO a , H 2 S0 3 , H 2 Si0 3 ). 
An amphoteric oxide is a metallic oxide which can show both basic 
and acidic properties, i.e., can react with both acid and alkali to 
produce a salt and water only. 
Examples are ZnO and A1 2 3 . 

Basic ZnO + H 2 SO« -»- ZnS0 4 + H 2 

Al a 8 + 6HC1 -► 2A1C1„ + 3H 2 
Acidic ZnO + 2NaOH + H 2 -*■ Na,Zn(OH) 4 

AJ.O, + 2NaOH + 3H s O -»• 2NaAl(OH)« 


The salts from the acidic reactions are sodium zincate and sodium 

A neutral oxide is an oxide which shows neither basic nor acidic 
character (as defined above), e.g., dinitrogen oxide, carbon mon- 


1. Under what conditions does oxygen react with the following sub- 
stances: (a) phosphorus; (b) sulphur; (c) magnesium; (d) copper? Name 
the products and mention any manner in which you could classify them. 

2. How would you employ pure nitrogen monoxide (contained in a 
vessel under pressure) to estimate the volume of oxygen in a sample of air 
enclosed over water in a graduated tube? 

3. Describe the preparation and properties of oxygen. If you had two 
vessels, one containing air and the other oxygen, how would you dis- 
tinguish them by simple tests? (O. and C.) 

4. Describe an accurate method of finding the percentage of oxygen in 

I (i,» Air 

How do you account for the percentage being so nearly constant when 
there are many causes which remove oxygen from the air? (L.) 

5. Manganese dioxide is said to catalyse the decomposition of potassiuni 
chlorate; what does this statement mean ? Describe experiments you would 
carry out to prove that the statement is true. (O. and C.) 

6. Describe one method for the quantitative estimation of oxygen in the 
air. State the precautions which are necessary to obtain an accurate result. 
Calculate the density of air (in terms of hydrogen) on the assumption that 

air is a mixture of four volumes of nitrogen and one volume of oxygen. 
(H = 1; N - 14; O - 16.) (N.U.J.B.) 

7. Outline the usual laboratory preparation of oxygen. How is it 
collected ? (Details of apparatus are not required.) How is oxygen obtained 
from the air? 

Describe an accurate method for measuring the percentage by volume ot 
oxygen in the air. (Methods involving the use of a bell-jar should not be 

State briefly how oxygen may be converted into ozone. (N.U.J.B.) 

8. Describe two experiments which you could perform in the laboratory, 
one in which oxygen justifies its name of acid-producer and one which 
gives an opposite result. What evidence is there that the gases of the 
atmosphere are not chemically combined with each other? 

9. Describe the usual preparation and collection of oxygen in the 
laboratory. State how oxygen may be converted into (a) an acidic oxide; 
(b) an alkaline oxide; (c) an insoluble basic oxide; (d) an allotropic form. 
Give also a sketch illustrating (</). (L.) 



Chapter 21 

Hydrogen Peroxide and Ozone 


THE commoner oxide of hydrogen is water, H 2 0. It forms, in 
addition, another oxide, hydrogen peroxide, H 2 2 . This com- 
pound is a colourless liquid and has been known for more than a 
century. Its purification is difficult and it was not until 1894 that it 
was obtained in a pure state. 

Preparation of hydrogen peroxide, H 2 O g 

Hydrogen peroxide may be prepared by acting upon the peroxides 
of certain metals with acids. The materials usually used are barium 
peroxide and dilute sulphuric acid, because the barium sulphate 
produced is insoluble and can be filtered off. 

BaO s + H 2 S0 4 -* BaS0 4 j + H 2 2 

To 200 cm 3 of water add 20 cm 3 of concentrated sulphuric acid. 
Place the beaker containing the dilute acid in a freezing-mixture of 
ice and salt and allow it to cool. Gradually add to it a quantity of 
previously moistened hydratcd barium peroxide until the mixture 
only just reacts acid. Then allow the mixture to settle and filter it. 
Add to the filtrate a few drops of baryta water (barium hydroxide 
solution) until it is accurately neutral, and the resulting aqueous 
solution of hydrogen peroxide is ready for use. 

The above method of making hydrogen peroxide gives a pure 
solution, but the yield is small unless great care is taken with the 

For the reactions of hydrogen peroxide make a solution as follows: 

Dilute about 5 cm 3 of syrupy phosphoric acid with its own volume 
of water. Cool under the tap and add in small portions at a time two 
or three saltspoonfuls of barium peroxide. Filter the solution and 
use it for the reactions described below. 

3BaO., + 2H 3 PO« -*■ Ba 3 (P0 4 ) 2 | + 3H 2 2 

Note. If the solution is acidic it will contain some barium ions but 
these do not interfere with the tests. 

Properties of hydrogen peroxide 

The pure compound is a syrupy liquid. It is usually used in dilute 
solution in water. 

Action of heat on hydrogen peroxide 

Warm hydrogen peroxide solution in a test-tube. Effervescence 
occurs. The gas given off is oxygen. The gas will not rekindle a glow- 
ing splint because of the presence of steam. 

2H 2 2 -*2H 2 + 2 

It will be readily understood from the above action that hydrogen 
peroxide acts as an oxidising agent. 

Action of hydrogen peroxide on lead(II) sulphide 

Precipitate lead(II) sulphide by passing hydrogen sulphide into a 
solution of lead(II) nitrate in a boiling-tube. Allow the precipitate to 
settle, pour off the liquid and add to the black lead(U) sulphide some 
hydrogen peroxide solution. Leave it to stand for some time, shaking 
occasionally. The precipitate gradually turns white, because it is 
slowly converted to lead(II) sulphate. 

PbS + 4H.O, ->■ PbS0 4 + 4H t O 
Hydrogen peroxide is here acting as an oxidising agent, oxidising 
lead(II) sulphide to lead(II) sulphate, and being itself reduced to 

This reaction is used in restoring pictures. Hydrogen sulphide in 
the air reacts with the white lead paint (lead(II) carbonate) of the 
picture to produce lead(II) sulphide, which is brown and makes the 
picture dingy. Washing with hydrogen peroxide restores the white 

Action of hydrogen peroxide on acidified potassium iodide solution 

Acidify a solution of potassium iodide with dilute sulphuric acid. 
Add hydrogen peroxide. A brown coloration is caused by the pro- 
duction of free iodine. 

2KI + H 2 S0 4 + H 2 2 -»■ K 2 S0 4 + I, + 2H,0 
Here again, hydrogen peroxide is an oxidising agent, oxidising 
potassium iodide to iodine, and being reduced to water. 
In ionic terms, this equation is: 

21- + 2H+ + H,O s -»- 2H,0 + I, 





The iodine ion is oxidised by electron loss, as: 2I - — 2e~— y I,. The 
electrons are accepted by the oxidising agent, as: 

2H+ + H a O a + 2c--> 2H,0 
That is, hydrogen peroxide is reduced by electron gain. 

Hydrogen peroxide also oxidises an iron(Il) salt, e.g., FeS0 4 , in 
acidic solution to the iron(III) state, the solution turning from green 
to yellow. 

2Fe* + + H a O a + 2H" -> 2Fe» + + 2H a O 
A solution of a soluble sulphite is similarly oxidised to sulphate. 

scy- + H a O a -*- so 4 «- + H t O 

Its powerful oxidising action makes hydrogen peroxide useful as a 
bleaching agent. It oxidises many dyes to colourless substances, 
without damaging the fabric carrying the dye. It bleaches hair to the 
well-known 'peroxide blonde' colour, and is also used for bleaching 
the more delicate materials, such as silk or feathers. It is also em- 
ployed for cleansing wounds. 

Hydrogen peroxide in reducing actions 

Action on kad(IV) oxide. Suspend some lead(TV) oxide in dilute 
nitric acid. Add hydrogen peroxide. Effervescence occurs with evolu- 
tion of oxygen (rekindles a glowing splint). A colourless solution 
remains. In this reaction, lead(lV) oxide is converted to Iead(II) 
oxide, PbO, which dissolves in the nitric acid. Hydrogen peroxide is 
converted to water. 

PbO, + H a O a -*• PbO + H a O + O a 
PbO -f 2HNO a -»- Pb(N0 3 ) a + H 2 

Similar reactions occur, with evolution of oxygen, when hydrogen 
peroxide reacts with silver oxide, acidified potassium permanganate 
solution and ozone. 

Ag a O + H a O a -v 2Ag -f- H a O + O t 
Silver oxide is reduced to metallic silver (a black precipitate). 
2Mn0 4 - + 5H a O a + 6H+ — > 2Mn 2+ + 8H a O + 50, 
The permanganate ion is reduced to a manganese salt. 

O s + H a O a -> H a 4 20 a 
Ozone is reduced to oxygen. 

In these reactions, hydrogen peroxide appears to act as a reducing 
agent by losing electrons (in association with hydrogen ions), as: 

H a O a -* 2H+ + O a + 2e~ 
The electrons are accepted by the other reacting substance, acting as 
an oxidising agent, as: 

(Ag + ) a O a ~ + 2H+ + 2e~ -> 2Ag + H a O 
MnO«~ + 8H+ + 5e~ -»- Mn a+ + 4H a 
PbO a + 2H + + 2e~ -> PbO + H a O 
Hydrogen peroxide is decomposed cataly tically by many substances, 
e.g., manganese dioxide, finely powdered gold and platinum. 
2H a O a ^-2H a O + O a 
Add a pinch of manganese dioxide to about 5 cm 3 of hydrogen 
peroxide solution in a test-tube. Oxygen is rapidly evolved. 

Sale of hydrogen peroxide 

Hydrogen peroxide is sold retail in *10 volume* and '20 volume' 
solutions, i.e., 1 cm 3 of the solution yields 10 cm 3 or 20 cm 3 of 
oxygen at s.t.p. when heated. To minimise loss by catalytic decom- 
position, the solutions should be as pure as possible and the con- 
tainers free from roughness. Some additives, e.g., acetanilide, can act 
as stabilisers (negative catalysts) but purity is the best safeguard. 


Ozone can be prepared from oxygen by the use of the silent electrical 
discharge. The apparatus is shown in Fig. 86. 

Dry oxygen is passed through the space between the glass tubes. 
Each tube is coated with tin-foil (T), which is connected to the 
terminals of an induction coil. No actual sparking takes place between 
the layers of tin-foil, nor does the oxygen come into contact with 
them, but a state of electrical strain exists. The issuing gas may con- 
tain up to 5% of ozone. This ozonised oxygen should not be allowed 
to come into contact with rubber, which is attacked by ozone. 

30 a -*■ 20 3 
Other forms of apparatus may be used for making ozonised oxygen. 

>. Ozonised 



Fig. 86. 
Preparation of ozone. 



Their principle is the same as above, but dilute sulphuric acid takes 
the place of the layers of tin-foil. 

Tests for ozone. It possesses a smell which resembles that 
of very dilute chlorine. (It is noticeable near an electrical 
machine in operation.) 

Ozone oxidises mercury and makes the mercury 'tail', i.e., 
leave a trail of mercury stuck to glass as the mercury flows 
across it. Quite small traces of ozone can be detected in this 

Properties of ozone, 3 

Ozone is a gas at ordinary temperature and pressure. It is obtained 
pure by liquefaction and is then a dark blue, explosive liquid, boiling 
at about — 1 12°C under ordinary pressures. 

Ozone as an oxidising agent 

Ozone is a vigorous oxidiser. It oxidises lead sulphide to lead 

PbS + 40 3 -»• PbS0 4 + 40, 
black while 

lead(II) lead(II) 

sulphide sulphate 

and hydrogen sulphide to sulphuric acid. 

H 2 S + 40 3 -»-H 2 S0 4 + 40, 

It also liberates iodine from potassium iodide in acidic solution 

2KI + H 2 S0 4 + O s -»- 1 ? + 2 + K t S0 4 -4- H 2 




This reaction is expressed in ionic terms as: 

21- + 2H + + O a -*■ I, + 2 + H 2 

The iodide ion is oxidised by electron loss: 2I _ — 2e~ —*■ I 2 ; the 
electrons are accepted by ozone acting as an oxidising agent: 

2H+ + O s + 2e--»- O, + H t O 
This powerful oxidising action of ozone makes it useful for ventila- 
ting places to which fresh air has little access. It attacks the organic 
compounds which are responsible for the 'stuffy' smell. Some of the 
London tube railways are ventilated in this way, metallic ozonisers 
of the type described above being used to supply ozonised air. The 
gas becomes poisonous at concentrations exceeding about 1 in 50 000 
of air, by volume. 



AHotropy of oxygen 

Oxygen and ozone are allot ropes (p. 287) of the same element; the 
difference between them is one of molecular complexity, oxygen 
having a diatomic molecule, 2 , and ozone a triatomic molecule, 3 . 
Their chemical identity is shown by the fact that each can be con- 
verted to the other (as shown in earlier pages) without change of 

The following table compares the two allotropes. 

Ozone, O, 

Oxygen, O, 

Gas at s.t.p. 
Density 16 (H = 1). 
Insoluble in turpentine. 
Heat has no action. 

No effect on mercury at room 

No effect on rubber. 
Has no effect on potassium iodide 


Oxidising agent. 

Gas at 

Density 24 (H = 1). 

Absorbed by turpentine. 

Heat converts ozone into oxy- 
gen. 20 3 -»■ 30 2 . 

Makes mercury 'wet' glass 

Attacks rubber. 

Liberates iodine from potas- 
sium iodide solution. 

(Equation opposite) 

Vigorous oxidising agent. 


1 How is ozonised oxygen prepared in the laboratory? Describe thrbb 
experiments to show how it diners from oxygen. If 100 cm" of pure ozone 
were heated and then reduced to the original temperature and pressure, 
what would be the volume of the resulting gas? (N.U.J.B.) 

2 How would you prepare a dilute solution of hydrogen peroxide? 
What is the action of hydrogen peroxide solution on (a) ozone; (/>) lead 
sulphide; (c) acidified potassium iodide solution? 

What is meant by a '10 volume' solution of hydrogen peroxide? 
3. How is ozonised oxygen obtained? 

In what respects docs this gas differ from ordinary oxygen? How has 
it been shown (a) that ozone is composed of oxygen atoms only ; (6) that 
its molecular weight is greater than that of oxygen? (O. and C.) 
4 (a) How would you prepare a dilute solution of hydrogen peroxide? 
How does hydrogen peroxide react with (i) ozone; (11) acidified 
potassium iodide solution? .„,-»,» j .«« 

(b) What volume of oxygen, measured at 136.5 C and 760 mm 
pressure, could be obtained by boiling 100 g of a solution con- 
taining 17% by weight of hydrogen peroxide? (N.U.J.B.) 



5. Describe the properties of ozone. 

On partly ozonising 100 cm 3 of oxygen a decrease of volume of 10 cm' 
resulted. What volume of ozone had been produced ? The resulting gas 
was treated with excess of a solution of potassium iodide when the follow- 
ing reaction took place: 

3 + 2KI + H.O -> 2KOH + I, + O, 
Calculate (a) what volume of gas would remain; (b) what weight of iodine 
would be liberated, assuming the volumes to have been measured at 
s.t.p. (C.) 

6. How can oxygen be converted into its allotropic form? Compare the 
properties of oxygen with those of its allotropic form. (O. and C.) 

Chapter 22 

Carbon and its Oxides; Flame 




PURE carbon is found in the form of diamond (India, South Africa) 
and impure carbon as graphite (Ceylon). Carbon is a constituent 
of numerous naturally occurring substances such as coal, mineral oils, 
carbonates, organic matter of all kinds and occurs in the air to a small 
but very important extent (0.03-0.04% by volume) as carbon dioxide 
(see p. 185). 

AUotropy of carbon 

If an element can exist (without changing its state) in two or more 
different forms, the element is said to exhibit allotropy. 

The forms of the element are known as allotropes of it. They always 
exhibit different physical properties and may have different chemical 
properties also. 

Carbon exists in the following allotropic modifications: 

Diamond. The diamond is in the form of octahedral crystals of 
density 3.5 g/cm s which have, when cut and polished, an amazing 
lustre which makes them valuable as jewellery. It is also the hardest 
substance known and has a commercial value for the manufacture of 
glass cutters and rock borers. 

Graphite. Graphite exists as black, slippery, hexagonal crystals. It 
is found naturally as plumbago and is manufactured artificially by 
heating coke to a very high temperature in the electric furnace 
(Acheson process). 

It is used extensively in the manufacture of lead pencils (which 
contain a long thin cylinder of graphite and clay) and also as 'black 
lead* as a protective coating for iron articles. It is used as a lubricant, 
particularly for small bearings (for example, those in dynamos and 
vacuum-cleaner motors) which require little, but regular, lubrication. 






Amorphous carbon (which is really made up of small crystals of 
graphite) exists in many forms: 

Animal charcoal \s made by heating animal refuse and bones with a 
limited supply of air. It also contains much calcium phosphate. It 
has the property of absorbing colouring matter (for example, litmus) 
and has a use in industry in removing the colouring matter from 
brown sugar. 

Wood charcoal is made by heating wood with a limited supply of 
air. It is a light porous variety, and is a remarkably good absorbent 
for gases (1 cm 3 of wood charcoal will absorb nearly 100 cm 3 of 
ammonia gas at 0°C). 

Lampblack is made by burning oils (for example, turpentine) with 
a limited supply of air, and it is used for making printers' ink and 

Sugar charcoal is a very pure form of carbon, and is made by 
removing the elements of water from sugar. 

Coke, gas carbon and soot are other forms of impure amorphous 

Crystal structure 

In diamond, covalency operates between the carbon atoms throughout 
and produces a crystal which is one single molecule which may become 
very large (a macromolecule). The crystal unit contains five atoms, ABCDE 



Dots are centres of carbon atoms 

Fig. 87. 
Crystal structure of diamond and graphite. 

(Fig. 87), and it is repeated indefinitely, forming interlacing hexagons. The 
strength and uniformity of bonding make diamond very hard, non-volatile 
and resistant to melting and to chemical attack. In graphite, carbon atoms 
arc combined by covalency in hexagons in parallel planes. Between the 
planes, bonds are much weaker and are the result of vander Waals forces 

(o 80) These weak forces allow movement of the planes parallel to each 
other and this makes graphite very soft. The open structure (Fig 87) makes 
cranhite more liable than diamond to chemical attack. (See also below.) 
In the three-dimensional formation of diamond, all four valency electrons 
oer atom are involved in covalent bond formation with four adjacent 
carbon atoms. In the parallel atomic layers of graphite, only three valency 
electrons per atom are definitely located in bond formation, the bonding 
between layers being van der Waals forces. Consequently, some electrons 
in graphite are mobile and allow it to conduct electricity, which diamond 
does not. 

Experimental evidence that graphite and diamond are allotropes of 
carbon is the following. Both these substances can be shown to burn 
when heated in excess of oxygen and the sole product in each case is 
carbon dioxide. Further, if the evidence is made quantitative by 
absorbing the carbon dioxide in potassium hydroxide solution or 
soda lime, it can be shown that the weight of carbon dioxide obtained 
from one gram of diamond is the same as from one gram of graphite, 

i.e., 3.67 g. 
The following table summarises a comparison ot diamond with 




Density 3.5 g/cm 3 average; vari- 

Colourless, transparent; very 
high refractive index 

Hardest known natural sub- 

Electrical non-conductor 

Not attacked by these reagents 

Density 2.3 g/cm 3 average; vari- 
Black, opaque 

Very soft, marks paper 

Good electrical conductor 
Attacked by potassium chlorate 
and nitric acid together 

Diamond is transparent to X-rays while glass is almost opaque. This 
test distinguishes the two. 

Properties of carbon 

Carbon is not a very reactive element. All forms of carbon can be 
made to burn in excess of oxygen to form carbon dioxide, although 
the temperature at which they commence to burn varies. As the 
carbon burns a great amount of heat is liberated. 

c + o 8 ->co. 

Sulphur will also combine with carbon at a high temperature to 

form carbon disulphide. 

C + 2S-»-CS 2 





Owing to the fact that carbon combines readily with oxygen, it acts 
as a reducing agent and is used in industrial practice in obtaining iron 
and zinc from their ores. (See p. 483, and p. 481.) 

Reducing property of carbon. Scrape a small hole in a charcoal 
block and place in the hole a mixture of lead(II) oxide and anhydrous 
sodium carbonate. (The carbonate melts and forms a protective coat- 
ing, preventing the metal from being oxidised.) Turn the luminous 
Bunsen flame low, direct a jet of flame by means of a mouth blowpipe 
on to the mixture and heat it for a few moments. Allow to cool and, 
on ejecting the substance from the hole, you will find a small grey 
globule of metallic lead which can be cut with a knife and which will 
mark paper. 

PbO-f C->-Pb + CO 

In this reaction, carbon in association with the O 2- ion of the oxide 
makes electrons available and, therefore, acts as a reducing agent, as : 

0«- + C-»-CO + 2e- 

The ion, Pb 2+ , of lead(II) oxide is reduced by accepting these elec- 
trons, as: 

Pb a+ + 2e~ -*■ Pb 

Carbon is insoluble in all common solvents, a fact which all motor- 
car owners know to their cost Petrol consists of hydrocarbons similar 
to methane, CH«. In the cylinders these hydrocarbons are oxidised 
by the oxygen of the air. If the supply of air is insufficient a deposit 
of carbon is left inside the cylinders as a hard black solid. This would, 
in time, choke up the engine and hence the carbon has to be removed. 
Since carbon is not soluble in any common solvent, this has to be 
done by dismantling the engine and removing the carbon mechanic- 


Coal is an impure form of carbon and, as such, is used as a domestic 
source of heat. Its use is, however, uneconomic, as a considerable 
amount of the heat of the combustion goes up the chimney and also 
many valuable products are lost. Further, much of the carbon escapes 
as soot, polluting the air in the larger industrial areas. 

The essentials of coal-gas manufacture in its 'classical' form (i.e., 
until recent changes have modified it) are the following. A non- 
caking, bituminous coal, finely powdered, is fed continuously down 
vertical fireclay retorts heated to about 1000°C by producer gas. The 
volatile products pass by the ascension pipe to the hydraulic main, 
where coal-tar and an aqueous liquor begin to condense. 

The gas passes to water-cooled condensers, where more tar and 
aqueous liquor separate out. All these liquids are fed to the tar-well, 

where coal-tar separates as a lower layer with gas-liquor as an aqueous 

layer above it. 

The gas then enters scrubbers, where it encounters a spray of water. 
This dissolves out any remaining ammonia and also acidic gases 
(HjsS, C0 2 , HCN) as their ammonium salts. This solution is added 
to the gas-liquor. The impurity, carbon disulphide, in the coal-gas 
is now converted to hydrogen sulphide by passage over nickel 
(catalyst) at 450°C. 

CS 2 + 2H 2 -*-2H 2 S-i-C 

AH the hydrogen sulphide of the gas is now removed in the purifiers, 
by passage over hydrated ferric oxide. 

Fe 2 O s + 3H t S -»• Fe»S 3 + 3H 2 
Fe 2 3 + 3H 2 S -»• 2FeS + S + 3H a O 
After use, the material is exposed to air to regenerate the oxide and 
is used again. 

2F ei S s + 30 2 ->2Fe 8 3 + 6S 
4FeS + 30 5 -*■ 2Fe.0 3 + 4S 
When the sulphur content reaches 55%, the spent oxide is sold as a 
source of sulphur dioxide (p. 407). 

The gas is then scrubbed with petroleum oil to remove benzene, 
toluene and naphthalene, which are recovered by fractional distilla- 
tion, after which it enters a gas-holder and is delivered to consumers. 
The white-hot coke remaining in the lowest third of the retort has 
steam passed through it. This produces water-gas, which mixes with 
the coal-gas and improves its thermal quality. 
C + H,0-»-CO + H, 
Sulphur (in spent oxide) is a by-product and hydrogen cyanide is 
recovered as sodium ferrocyanide. 

To summarise: 

distilled at about I000°C 


hydrogen, 50%, 
methane, 30%, 
carbon monoxide, 
8% and small 
amounts of 
ethylene, CO t 
and other gases. 



! is distilled with | is distilled for 
lime. Ammonia is benzene, phenol, 

\ evolved and ab- | toluene, cresol, 

' sorbed by dil. | pilch. 
H.SO, to yield 
ammonium sul- 

'■ phate (fertiliser). 


is used as a fuel, 
and as a source of 
water-gas and 
producer gas, and 
as a large-scale 
reducing agent. 



Petroleum is now being applied to the production of a fuel gas to 
replace or supplement coal-gas. 

Fuel gas from petroleum 

Naphtha (p. 324) is vaporised and purified from sulphur com- 
pounds which would poison the catalyst in the main reaction. The 
purified naphtha vapour and steam are passed, in two stages, at 750°C 
and 690°C, and at 28 atm. pressure, over a nickel catalyst. Taking 
C 5 H 12 as typical of naphtha, the reactions can be summarised as 

C 6 H It + 5H 2 0-»- 5CO + HH a 
CO + 3H„ ^ CH« + H s O 
CO + H t O^ CO a + H, 

The product is passed over heated iron oxide (catalyst) to reduce the 
carbon monoxide content. After cooling, some carbon dioxide is re- 
moved (by water under pressure or alkali) and the gas is dried. It has 
the approximate volume composition : carbon monoxide, 3 % ; methane, 
CH 4 , 34%; hydrogen, 48%; carbon dioxide, 15%. Because of the low 
carbon monoxide content, this gas is much less poisonous than coal- 
gas, and it satisfies British standards of calorific value. A plant using 
this process is said to need only one-eighth of the space and one-sixth 
of the capital cost of a corresponding coal-gas plant. 

Recent borings in the North Sea have revealed reserves of natural 
gas (mainly methane) estimated (May 1969) to yield 3000 million 
cubic feet of gas per day for. perhaps, 35 years. This daily output 
would meet the whole of Britain's current needs of gaseous fuel. Pipe- 
lines will have to be laid and appliances modified to correspond with 
the fact that natural gas has twice the calorific power of present gas 
supplies (volume for volume). Methane is also being imported, in 
refrigerated tankers as liquid, from the Saharan oil-fields. 

Smokeless fuels 

The objections (p. 290) to the burning of 'raw* coal have led to the 
investigation of smokeless fuels. Coal-gas is, of course, smokeless and 
so is coke. But coke is hard to ignite and burns badly in the ordinary 
domestic grate. 

Low-temperature carbonisation of coal has been tried. Coal is dis- 
tilled in retorts as in making coal-gas, but the temperature of distilla- 
tion is much reduced (under 750°C instead of about 1000°C). Some 
coal-gas, oil and tar are produced. The solid residue is a fuel which 
ignites and burns much more readily than coke and is smokeless. So 
far, however, this solid fuel is too expensive compared with 'raw' 




Producer gas 

Producer gas is a fuel gas made by passing air through a thick layer 
of white-hot coke in a producer (Fig. 88). In the lower part of the 
producer, with air in excess, the reaction is: 

C -f 2 — >■ C0 2 (strongly exothermic; i.e., heat liberated) 
As the carbon dioxide rises through the mass of white-hot coke, the 
further reaction is: 

CO, -f C -> 2CO (endothermic; heat absorbed) 



Sheet iron with 
fire-brick lining 

Fig. 88. 
Manufacture of producer gas. 

On the balance of these reactions, net heat is liberated and the reaction 
can continue indefinitely. 

The gas emerging from the producer consists of carbon monoxide 
(one-third by volume) and unchanged nitrogen of the air (two-thirds). 
This producer gas is usually distributed, while still hot from the pro- 
ducer, to points on the site where heating is required; it is mixed with 
secondary air and burnt to produce more heat. 

2CO + O t — ► 2COj, (exothermic; heat liberated) 
Since two-thirds (by volume) of the gas is non-combustible nitrogen, 
its calorific value is not high and the gas is not distributed to con- 
sumers as town-gas is. It has been found very useful, however, in such 
processes as the firing of retorts and glass furnaces. 



Water-gas is produced by the passage of steam through a mass of 
coke at a temperature (minimum 1000°C) which allows the reaction: 

C -f H,0 — *■ CO + H, (endothermic; heat absorbed) 
giving equal volumes of carbon monoxide and hydrogen, both of 
which are combustible, so the gas has a high calorific value. 

As shown, the above reaction is endothermic, so the temperature 
of the coke falls. Below 1000°C, much steam passes without reaction 
and, also, incombustible carbon dioxide is formed. 

C + 2H 2 0-»-CO, + 2H 2 
To meet this situation, the plant similar to the producer above, is 
used in the following way. An air-blow of about two minutes raises 
the temperature of the coke to incandescence (perhaps 1400°- 
1500°C), using the exothermic reaction of producer gas (above). Then 
a steam-blow of about four minutes produces water-gas, which is 
collected. This cools the coke and the air-blow is renewed and so on 

Water-gas can be burnt as a fuel gas with air, both the following 
reactions being exothermic. 

2H, + O a -► 2H s O; 2CO + 2 -*■ 2CO a 

Water-gas has been an added constituent of town-gas for many years, 
being made by the action of steam on the white-hot coke residue in 
the retorts. The high carbon monoxide content makes the gas more 
poisonous. It has also been an industrial source of hydrogen and of 
organic chemicals, e.g., to make methanol. Water-gas and hydrogen 
are passed, with heating to 450°C, over a catalyst (oxides of zinc and 
chromium) at 200 atm pressure. 

CO + H 2 + H, -»- CH,OH 



This gas was first observed by Van Helmont towards the end of the 
sixteenth century, but Black (1728-99) first showed that the gas could 
be prepared by the action of dilute acids on calcium carbonate. It 
occurs in the air on the earth's surface to the extent of about 0.03% 
of its volume; it issues from rocks in volcanic regions, and occurs in 
mines as 'choke damp'. Certain mineral springs contain the gas and 
it is always present in natural drinking water because of its solubility 
in water. Its biological importance is dealt with on p. 185. 



Preparation or carbon dioxide 

It is usually made in the laboratory by the action of dilute hydro- 
chloric acid on marble. 

Place several pieces of marble in a flask (or bottle), as shown in 
Fig 89, and pour some dilute hydrochloric acid down the thistle 
funnel on to the marble. There is effervescence and a colourless gas 
is liberated which is collected over water 1 or by downward delivery, 
the gas being denser than air. Calcium chloride solution is left in the 

CaC0 3 (c) + 2HCl(aq) ->- CaCl s (aq) + H s O + C0 2 (g) 
or CO/" -I- 2H- — H 2 + C0 2 

If the gas is required pure and dry it can be passed through potas- 
sium hydrogen carbonate solution in a wash-bottle (to remove sus- 
pended hydrochloric acid), dried by passing it through a calcium 
chloride U-tube and collected by downward delivery (the gas being 
denser than air). 

Dilute Hydrochloric 




Fig. 89. 
Preparation of carbon dioxide. 

Collect several gas-jars of the gas, and perform the following experi- 

Effect of carbon dioxide on a lighted splint. Plunge a lighted 
splint into a gas-jar of the gas. It is extinguished. Carbon 
dioxide docs not support combustion. 

Action of carbon dioxide on lime-water. Pour lime-water 
into a gas-jar full of carbon dioxide. The lime-water goes 
milky. If the mixture is allowed to stand, you will see white 
solid particles separate out. These are particles of chalk. 

» There is some loss due to the solubility of the gas in water. Water, at ordinary 
temperature and pressure, absorbs its own volume of carbon dioxide. 

H 2 



The milkiness is due to a suspension of the insoluble sub- 
stance, chalk, in water. 

Ca(OH), + CO s -> CaCO a 

calcium carbon calcium 

hydroxide dioxide carbonate 

solution (chalk) 


The above test serves to distinguish carbon dioxide from 
any other gas. 

Effect of carbon dioxide on a lighted candle 

Lower a candle on a deflagrating spoon into a gas-jar of air. The 
candle can be extinguished by 'pouring' carbon dioxide into the gas- 
jar in which the candle is burning (Fig. 90). This shows carbon 
dioxide to be denser than air (density 22 relative to hydrogen). 

Fig. 90. 
Pouring carbon dioxide on to a candle. 

Effect of burning magnesium on carbon dioxide 

Lower a piece of burning magnesium into carbon dioxide in a 
gas-jar. It continues to burn for a short time with a spluttering flame, 
and black specks of carbon can be seen on the sides of the gas-jar. 
The- magnesium burns to magnesium oxide 

2Mg + CO i -*-2MgO- r -C 

This clearly shows carbon dioxide to contain carbon and oxygen. 

Solution of the gas in water 

Invert a gas-jar of carbon dioxide in a trough of cold water and 
shake the gas-jar. The water rises slowly, showing that the gas is 



soluble. Put a glass plate over the mouth of the jar and remove it. 
To the liquid in it, add blue litmus solution and shake. The solution 
becomes claret-coloured but not red. This is because carbon dioxide 
reacts with water to produce carbonic acid, which is, however, too 
weak to turn litmus solution red. 

HoO + CO, ^ H s CO s ^ 2H+ + CCV" 

watcr carbon carbonic (very slight) 

dioxide acid 

Carbon dioxide is an acidic oxide. 

Action of carbon dioxide with sodium hydroxide solution 

Repeat the above experiment, using sodium hydroxide solution 
instead of water. The rapid rise of the solution shows that the gas is 
quickly absorbed. The acidic carbon dioxide reacts with the alkaline 
solution producing sodium carbonate. 

CO g + 2NaOH -* Na 2 C0 3 + H,0 
or CO* + 20H- -* C0 3 *- + H 2 

This reaction is discussed more fully later (p. 312). 

Uses of carbon dioxide 

Solutions of the gas in water have a pleasant taste (the taste of soda 
water), and hence the gas is used in the manufacture of the effervescing 
drinks' called 'mineral waters*. The effervescence is caused by dis- 
solving the gas in water at a pressure of several atmospheres; when 
the pressure is released (by opening the bottle) the gas is liberated. 

Its use in the solid form is increasing. Carbon dioxide can be made 
into a white solid (carbon dioxide snow) by allowing liquid carbon 
dioxide to evaporate, the temperature falling to -78°C as the solid 
forms. This solid evaporates when heated, leaving no residue, and it 
is, therefore, used as a refrigerating agent for perishable goods. 

Owing to its non-inflammable nature, carbon dioxide is used for 
extinguishing fires. The usual fire extinguisher contains a solution of 
sodium carbonate which can be made to come into contact with 
dilute sulphuric acid by striking a knob. The carbon dioxide liberated 
forces a stream of effervescing liquid on to the fire, and the carbon 
dioxide prevents the air from getting to the burning material and so 
helps to put out the fire. For solid CO, crystals, see p. 73. 

Carbon tetrachloride, CCI 4 , tetrachloromethane 

This compound is made by chlorinating carbon disulphide, boiling 
under reflux, with iodine as catalyst. 

cs, + 3a,->cci4+s^a, 



The products can be separated by fractional 
distillation (boiling-points, CCI 4 77°C, S 2 CI 2 
138°C at standard pressure). The carbon tetra- 
chloride can be purified by shaking with dilute 
sodium hydroxide solution (to remove any 
chlorine) and then with water. This is done in a 
separating funnel (Fig. 91). Carbon tetrachloride 
is almost non-miscible with water and separates 
as the lower layer (density, 1.63 gem" 3 ) which 
is run off through the tap. The washing liquid 
(upper layer) is poured out through the top of the 
funnel. Anhydrous calcium chloride (bean size) 
is then added to the carbon tetrachloride and the 
mixture is left in a stoppered flask for at least 12 
hours. This will dry the carbon tetrachloride as 
the calcium chloride takes up the water present 
as the solid hydrate, CaCI a .6H 2 0. The liquid 
should then be decanted (or filtered through a 
dry filter-paper) into a dry distillation flask and 
distilled (apparatus as p. 191). If pure, the entire 
liquid should distil over at its boiling-point, 
77°C, at standard pressure. 
Properties and uses 

1. Carbon tetrachloride is a good solvent for fats and grease and 
is used as a de-greasing and dry-cleaning agent. It has the advantage 
of being non-inflammable. 

2. As pyrene, it is used as a fire extinguisher over small areas, e.g., 
a motor-car engine. Its vapour is very dense (about 5.5 times as dense 
as air) and tends to blanket the fire and extinguish it by excluding air. 
A disadvantage is that air at a very hot surface can oxidise it to 
carbonyl chloride, COCI 2 , a very poisonous compound. 

3. By heating with iron and water, carbon tetrachloride is con- 
verted to chloroform. 

CC1 4 + Fe + H + -> Fe 2+ + CI" + CHC1 3 
Otherwise, it is chemically rather inert; unlike most non-metallic 
chlorides, it is unaffected by water. 

The carbon tetrachloride molecule is shaped like the methane 
molecule (p. 75), substituting four chlorine atoms for the four 
hydrogen atoms of methane. 

Fio. 91. 
Separating funnel. 


Carbon monoxide is a poisonous, colourless gas with practically no 
smell. It is present in coal-gas and other gaseous fuels. It is formed by 



the partial combustion of carbon, and poisoning by the exhaust 
fumes of a motor-car in an enclosed space, for example a garage, is 
due to the presence of carbon monoxide. The blood of a person 
poisoned by the gas is a characteristic cherry-red in colour. An atmos- 
phere containing as little as 0.5% carbon monoxide may cause death 
if breathed for some time, and an atmosphere containing 0.1 / is 

Preparation from oxalic acid 

Fit up the apparatus as shown in Fig. 88. Place some oxalic acid 
crystals (H i C l O«.2H,0) in the strong flat-bottomed flask and pour 
concentrated sulphuric acid down the thistle funnel. Warm the mix- 
ture gently (always have the greatest respect for hot concentrated 
sulphuric acid). The white crystals dissolve, effervescence is observed, 
and a mixture of carbon monoxide and carbon dioxide gases is 
evolved. By passing the mixture through a concentrated solution of 
caustic potash the carbon dioxide is absorbed, and the carbon mon- 
oxide passes on and is collected over water in which it is insoluble. 
2KOH + CO,->- K 2 C0 3 + H a O 

Chemistry of the action. Oxalic acid has the formula H 2 C 2 0« and 
the hot concentrated sulphuric acid removes the elements of water 
from the molecule of oxalic acid, leaving a mixture of equal volumes 

oxalic acid crystals 

and concentrated 
I sulphuric acid 

Cone, potassium 

hydroxide soln. 

to absorb carbonyl carbon 
dioxide . 'monoxide 


Fio. 92. 

Preparation of carbon monoxide. 

of carbon monoxide and carbon dioxide. It is because the carbon 
dioxide is there in quantity, and not merely as a small trace of im- 
purity, that it is necessary to pass the gas through two wash-bottles 
containing potassium hydroxide solution. 

H 2 C s 4 -H 2 0-»-CO + C0 2 
oxalic water carbon carbon 

acid (removed monoxide dioxide 

by acid) 





Preparation from formic acid 

Fit up the apparatus shown in Fig. 93. Place one or two teaspoonfuls 
of sodium formate in the fiat-bottomed flask and allow concentrated 
sulphuric acid to run in from the tap funnel. The reaction takes place 
in the cold, effervescence is observed and carbon monoxide is col- 
lected over water in which it is insoluble. 





Fig. 93. 
Preparation of carbon monoxide. 

Chemistry of the action. Sodium formate is converted into formic 
acid from which the elements of water are immediately removed by 
the concentrated sulphuric acid : 

HCOONa + H 2 S0 4 -»- HCOOH + NaHS0 4 

by acid 

The above actions take place simultaneously and can be represented 
by the equation : 

HCOONa + H,S0 4 ->• NaHS0 4 + CO + H s O 

Test. Carbon monoxide burns in air with a blue flame, 
forming carbon dioxide (the latter will turn lime-water 
turbid, forming a precipitate of chalk). 

Collect a gas-jar of the gas and perform the following experiment: 
Action of lime-water. Take away the glass plate, quickly pour in a 
little lime-water and shake. There should be no turbidity. (It may be 
very difficult to get rid of all the carbon dioxide present. If necessary, 
shake first with potassium hydroxide solution.) Apply a light to the 
gas and it will burn with a blue flame, and on shaking, the lime-water 
will go turbid. 

2CO + O t - 

carbon oxygen 
monoxide (air) 

•2C0 2 

Reducing properties of carbon monoxide 

Take a hard-glass tube containing a porcelain boat full of lead(H) 
oxide and pass carbon monoxide through it (Fig. 94). After allowing 
a few moments for the hard-glass tube to be filled with the gas, light 
the jet and heat the lead(II) oxide. A glow will spread through the 
lead(Il) oxide and, on allowing the apparatus to cool, grey metallic 
lead will be observed. The carbon monoxide has reduced the lead(II) 
oxide to lead, being itself oxidised to carbon dioxide. 
PbO(c) + CO(g) -»- Pb(c) + CO,(g) 
The gas will similarly reduce copper(II) oxide and iron(IH) oxide 
to the metal. 

CuO + CO ->- Cu + COj 
Fe 2 O s + 3CO ->■ 2Fe + 3CO, 
In all these reactions, carbon monoxide, in association with the 
s - ion of the oxides, makes electrons available and so acts as a 
reducing agent, as: 

O*- + CO -»- C0 2 + 2e~ 
The metallic ion is reduced by accepting these electrons: 

Cu* + + 2e~ ->- Cu 

Pb ,+ + 2e~ -*■ Pb 

2Fe 3+ + 6e~ -»- 2Fe 

carbon monoxide burning 
to carbon dioxide 


lead (n) 

Fig. 94. 
Reduction oflead(II) oxide by carbon monoxide. 

Preparation of carbon monoxide from carbon dioxide 

Carbon monoxide may be prepared by passing a stream of carbon 
dioxide through an iron tube packed with pieces of carbon heated to 








Fio. 95. 

Preparation of carbon monoxide by passing carbon 
dioxide over red-hoi carbon. 

red heat (Fig. 95). Any carbon dioxide present is then absorbed by 
caustic potash solution and the gas collected over water. 

C + C0 2 ->2CO 

carbon carbon carbon 
dioxide monoxide 

Reactions in a domestic fire 

The reactions in a deep, brightly glowing coke or coal fire are 
similar to that by which carbon monoxide is produced in the last 
experiment. They are sufficiently explained in Fig. 96. 

At A plenty of air available. Carbon 
burns to carbon dioxide. 

C + O, -»• CO, 

At B ascending carbon dioxide is 
reduced by red-hot carbon to carbon 

CO, + C -*■ 2CO 
At the surface the hot carbon mon- 
oxide burns in the air (to form carbon 
dioxide) with a flickering blue flame. 
2CO+ 0,->2CO, 

Bars i 
graU\ \ »— ZfAir 


Fio. 96. 

The household fire. 




Definition. Flame is a region of combining gases which radiate light 

and heat. 

Examples of flame with which you are familiar are the burning in 
air of such materials as hydrogen, coal-gas, wood or coal. In the 
last case, wood or coal when heated give out gases which will burn 
in air and since the wood or coal are supplying these combustible 
gases, the flame is closely associated with the solid matter. In every 
flame, there is produced by the chemical reaction so much heat that 
the system is raised to incandescence. We know, on the other hand, 
many cases where gases combine slowly with little rise of tem- 
perature and no flame. An example of this type of reaction is the 
combination of hydrogen and chlorine when left in contact with 
each other in diffused daylight. 

The terms 'combustible material' and 'supporter of combustion 
are interchangeable. 

When considering a Bunsen flame in air, we usually refer to the 
gas coming from the burner as the 'combustible material' and the 
air as the 'supporter of combustion'. This is mainly because we live 

in an atmosphere of air (which contains oxygen, a very reactive 

gas) which surrounds any burning material. If, by accident, this 

atmosphere of ours were to consist of coal-gas, it would be possible 

to produce the flame of a Bunsen burner by 

pumping air down the pipe usually connected 

to the gas-works. To illustrate this, set up the 

apparatus shown in Fig. 97, which consists of 

a lamp chimney with a square of asbestos 

(containing a round hole) resting on the top 

of the chimney. The cork at the base is fitted 

with a glass tube for a gas inlet and a mica or 

metal tube at B. The gas is turned on, and the 

aperture at A is closed momentarily whilst the 

gas issuing from B is ignited. The aperture at 

A is now opened and the gas is lighted. On 

inspection, it will be seen that there is now a 

flame at the top of the mica tube which consists 

of air burning in an atmosphere of coal-gas. 

Structure of flame 

The structure of a flame varies according 
to the chemical composition of the gas which 
is burning. The general conical shape is brought 
about by several factors, the more important 

Fio. 97. 

Air burning in 




of which are the effects of convection and the necessity for further 
supplies of gas to search for air with which to burn. The gas burning 
immediately on issue from the burner, uses up the air in that region 
and the gas following has to seek its supply of air from more distant 
sources. At the same time the convection currents set up confine this 
search to an upward direction. The blue zone seen at the base of a 
Bunsen or candle flame (see Fig. 99) is caused by the upward stream 
of air impinging on the base of the issuing cone of gas. The inner 
zone near the burner contains unburnt gas, since the outside of the 
flame has obviously the best opportunity to come in contact with 
the air, and the inside of the flame, the least. 
We will consider the structure of some well-known flames: 
(a) The flame of hydrogen burning in air. This flame is, as we should 
expect, a simple one. Two zones only are produced, a zone of unburnt 
gas surrounded by a zone in which combination of oxygen and hydro- 
gen takes place. The flame is almost invisible in dust-free air (See 
Fig. 98). 

of comoi nation 
'nburnt yua 


Fio. 98. 
Hydrogen flame. 

(b) The candle flame. Candle-wax consists of hydrocarbons, both 
the carbon and the hydrogen being combustible materials. It is as- 
sumed that both elements do not burn completely to form water and 
carbon dioxide except at the very outside of the flame where there 
is plenty of air. In the bright yellow zone there is incomplete com- 
bustion and particles of carbon raised to a white heat are present in 
it to give the flame most of its luminosity, whereas the outside zone 
where combustion becomes complete is only faintly visible. The zone 
round the wick consists of unburnt gas, and the blue zone at the base 
is a zone of rapid burning caused by the upward rush of air (due to 
convection) first meeting the combustible gases. (See Fig. 99.) 

The luminous Bunsen flame (air holes closed) gives a parallel with 
the candle flame, the gases in the coal-gas being methane (a hydro- 
carbon), hydrogen and carbon monoxide. 

(c) The Bunsen flame. The flame of a burning hydrocarbon or the 
luminous Bunsen flame is not a very hot flame and deposits soot on a 
cold article held in it. Bunsen devised a simple burner (Fig. 100) to 
ensure more complete combustion by introducing a supply of air 
entering with the gas, so that this supply of air together with the 




Candle flame. 

Luminous coal-gas flame. 
Fig. 99. 

| Stream of gas draws m 
J/ "^iair from outside when 
oirWes open 


Fio. 100. 

Bunsen burner. 



external supply is sufficient to produce complete combustion. Hence 
the flame with the air holes open is more compact, much hotter and 
non-luminous. The structure of this flame is shown in Fig. 101, and it 
will be seen that the luminous zone has been replaced by a zone in 
which the coal-gas burns with the internal supply of air. There is a 
limit to the amount of air which can be supplied by the holes at the 
base, for if sufficient is introduced to cause almost complete com- 
bustion, the rate of burning of the mixture exceeds the speed at which 
the gas is moving up the tube and the flame 'strikes back*. 

Gas burning trffn air 
mainly supplied 
from outside 

Gas burning with air 
supplied from base 

Unburnt gas 

Fio. 101. 
Non-luminous Hansen flame. 

Experiments to illustrate the structure of flame. The existence of a 
zone of unburnt gas and zones of varying temperatures may be 
shown by exploring the flame in the following way: hold a piece of 
paper or cardboard horizontally in the flame and remove it just be- 
fore burning would occur (Fig. 102). The inner zone will be found to 
cause no burning of the paper, whilst the hot parts of the flame will 
turn the paper brown. Alternatively, copper gauze may be held in 
the flame by means of tongs. 

Humphreys and Glasgow Lid. 

Plate 5. (a) (Above): Manufacture of 
carburcucd water-gas. Producer-gas 
and water-gas arc made in turn in the 
tapered vessel (/<•//). Oil is introduced 
in the two large cylinders to improve 
the calorific value of the gas. 

Plate 5. (b) (/.<//): Ammonia con- 
venors in the I.C.I, works at Billingham. 

r— ■Jul f*i<*„.i.„i iii.Unii', i /././. 


Plate 6. Manufacture and use of fertiliser. Fison's nitric acid plant at Stanford-le-Hope 
and (below) combine drilling corn and fertiliser. 


FLAME 307 

The causes of luminosity of flame 

The luminosity of a flame is affected by alteration of temperature 
and pressure of the burning gases and also by the presence or absence 
of solid particles. 

Fio. 102. 
Exploring temperature of flame. 

The presence of solid particles in a flame increases luminosity. 
Sprinkle a few iron filings into the non-luminous Bunsen flame. 
Sparks are formed as each particle is raised to a white heat. Simi- 
larly, platinum foil or a porcelain rod become white hot and emit 
light when heated to a high temperature in the non-luminous Bunsen 
flame. This principle is used in the incandescent mantle where small 
particles of thorium and cerium oxides are suspended in a non- 
luminous Bunsen flame. 

Hold a cold evaporating dish by means of tongs for a minute in the 
luminous Bunsen flame. On removal, the dish will be found to be 
blackened by carbon. No such effect is observed if the experiment is 
repeated using the non-luminous flame. The presence of the particles 
of carbon is thought to be responsible for the luminosity of this type 
of flame. 

Effect of increase of pressure. Increase of pressure of the gases 
taking part in the combustion increases the luminosity of the flame. 

The effect of temperature on luminosity. Set up the apparatus of 
Fig. 103 with a silica tube extension fitting over aBunsenburnertube. 
Open the air holes and obtain the non-luminous flame at the end of 



the silica tube. Now heat the silica tube strongly with a second 
Bunsen burner. As the silica tube is heated, the flame at the end of the 
silica tube gradually becomes luminous and luminosity diminishes as 
the tube is allowed to cool. 

Fig. 103. 
Effect of temperature on luminosity. 


Explosions result from very rapid, exothermic chemical reactions. 
Consider a mixture of hydrogen, two volumes, and oxygen, one 
volume. If the mixture is sparked, the gases are heated to very high 
temperature near the spark. They combine and liberate heat. 
H, + iO, -> H t O; A// - -286 kJ g-eqn" 1 

This heat raises the temperature of neighbouring gas, which com- 
bines liberating more heat, and so on with great rapidity. The mass 
of gas is raised to incandescence in a fraction of a second and the 
consequent great expansion produces a pressure wave in the air. The 
whole effect is called an explosion. 

Questions on this chapter will be found on page 317. 

Chapter 23 

Carbonates and Hydrogen Carbonates 











of these 
in water. 


Carbonates of these metals' 
not decomposed by heat. 

Carbonates of these metals 
decomposed into oxide of 
the metal by heat 

(Al forms no carbonate) 

Any carbonate 
with any acid 
liberates car- 
bon dioxide. 

Ammonium carbonate is also soluble in water. 

The above table summarises briefly the important properties of the 
common carbonates. 

Test for any carbonate 

Put some of the suspected carbonate in a test-tube and add dilute 
nitric acid. If a carbonate is present there will be effervescence and the 
gas which comes off will turn lime-water milky: 
OV" + 2H + -*■ H g O 4- CO, 

Ammonium carbonate (NH 4 ) 2 C0 3 

This compound is prepared as a sublimate, by heating ammonium 
sulphate with limestone. 

(NH 4 ) 2 S0 4 + CaCO s ->- (NH 4 ) s CO, + CaS0 4 
It is used as a constituent of 'smelling salts' and as a chemical reagent. 

Potassium carbonate, K 8 C0 3 

Potassium carbonate is very similar to sodium carbonate. (See 
p. 3 10.) It cannot, however, be made by the Solvay process because 




the hydrogen carbonate of potassium is too soluble. It is made by the 
Leblanc process and in other ways. 

Potassium carbonate differs chiefly from sodium carbonate in that 
it is very deliquescent, and is anhydrous. 

It is used to make soft soap, hard glass and potassium salts gener- 

Sodium carbonate (soda ash), \a.,C O , 

Sodium carbonate is obtained in both the anhydrous and crystal- 
line states by the Solvay process. 

Solvay process 

Very concentrated brine (28% NaCl) is saturated with ammonia 
gas in a tower and the ammoniacal brine is run down further Solvay 
towers up which carbon dioxide is forced. The towers are fitted with 
perforated mushroom-shaped baffles at intervals (Fig. 104). These 

Waste 4|| brine 


Suspension of 
sodium hydrogen 









Fio. 104. 
Solvay tower. 

baffles delay the flow of liquid and present surfaces for reaction. 
Sodium hydrogen carbonate is formed. It is not very soluble in 
water, so it precipitates; precipitation is assisted by cooling the 
lowest third of the tower. 

NaCl + NH 4 OH + CO a -»• NaHC0 3 + NH 4 CI 



Sodium hydrogen carbonate is filtered from the white sludge and 
washed free from ammonium compounds. It is then heated to convert 
it to sodium carbonate and carbon dioxide evolved is used again. 
2NaHC0 3 -v Na 2 CO s + H a O + CO» 

The substance formed is anhydrous sodium carbonate, which finds 
a wide market. If the crystalline form (washing soda) is required, the 
anhydrous solid is dissolved in such an amount of hot water that 
crystallisation occurs on cooling. The crystals are removed and 
allowed to dry. 

Na a CO s I- lOH^O-*- Na 8 CO 3 .10H 2 O 

Efficiency of the process 

(1) The principal raw materials for the process are cheap and 
plentiful. They are sodium chloride and limestone. The common salt 
is extracted from deposits as brine; the limestone yields quicklime 
and carbon dioxide when heated. 

CaC0 3 ?± CaO + C0 2 
Carbon dioxide is passed to the carbonating tower and approximately 
half of it is recovered in the heating of sodium hydrogen carbonate 

Ammonium chloride, NH/Cl", is left in the solution after pre- 
cipitation of sodium hydrogen carbonate or is washed out of the 
precipitate. Ammonia is recovered by heating the solution and wash- 
ings with quicklime (above) and returned to the ammonialing tower. 

CaO + H 2 -»- Ca(OH) 2 
2NH 4 C1 + Ca(OH), -> CaCl 2 + 2H 2 + 2NH 3 
Consequently, once the pipe-line is full, no further external supplies 
of ammonia arc theoretically required. In practice, about 2% of the 
ammonium chloride in use is added per circuit to restore manipu- 
lative losses. 

(2) The process is one of continuous flow and only a minimum 

labour cost. 

(3) The principal weakness of the process is its failure to utilise 
the chlorine of the sodium chloride used. This is lost as calcium 
chloride (see equation in (1) above) for which there is only slight 

Uses of sodium carbonate 

Four of the important uses of sodium carbonate are the following. 

(a) Manufacture of glass. Ordinary bottle glass is made by fusing 
together sodium carbonate (or sulphate), calcium carbonate, silica 





and a little carbon (reducing agent). Broken glass (culiet) is added to 
assist fusion. 

Na,C0 3 + SiO, -> Na 2 Si0 3 + CO, 
CaCO a + SiO, — »• CaSiO a + CO, 

The mixture of silicates (with some unchanged silica) constitutes the 

(b) Manufacture of water-glass. Sodium carbonate is fused with 
silica to produce sodium silicate. 

Na,C0 3 + SiO, -► Na,Si0 3 + CO, 

On cooling, a glassy solid is left, which is broken up, boiled with water 
and evaporated to a colourless, treacly material, known as water- 
glass. It is used in preserving eggs, in fire-proofing, and in producing 

(c) In domestic water-softening. Calcium ion, Ca 2+ , which is the 
principal cause of hardness in water, is precipitated from the water 
as chalk, Ca 2+ C0 3 2_ , by the addition of sodium carbonate. 

(d) In the manufacture of sodium hydroxide (Gossage's process). 
A 10% solution of washing soda is steam-heated and stirred in an 
iron tank, this metal having no action with sodium hydroxide (the 
product required). Lumps of quicklime are contained in steel-wire 
cages which dip into the liquid. 

CaO + H,0 -> Ca(OH), 
Na,CO s + Ca(OH), -»• 2NaOH + CaC0 3 

The second reaction is, in fact, somewhat reversible, but the lower 
solubility of calcium carbonate (compared with calcium hydroxide) 
causes it to precipitate and so force the reaction to the right (as writ- 
ten above). This produces a satisfactory yield of sodium hydroxide. 
Calcium carbonate is filtered off under reduced pressure and the 
filtrate is concentrated by steam heat, also under reduced pressure. 
Solid sodium hydroxide is not produced; the concentrated solution 
is applied to the required uses. For a laboratory version of this 
process, see p. 255. 

Laboratory preparation of sodium carbonate 

(a) By the action of heat on sodium hydrogen carbonate (see p. 124). 

(b) From sodium hydroxide (caustic soda) solution. Fit up the 
apparatus as shown (Fig. 105). Pass carbon dioxide (free from hydro- 
chloric acid) into a moderately concentrated solution of sodium 
hydroxide for some time until finally a white solid (sodium hydrogen 
carbonate) appears on the bottom of the boiling-tube. 

2NaOH + CO, ->■ Na,CO s + H,0 (1st stage) 
Na,C0 3 + H,0 + CO, -»■ 2NaHC0 3 (2nd stage) 

In ionic terms: 

20H- + CO, 
CO.*-+ HjO + CO, 

C0 3 *"4 H,0 (1st stage) 
2HC0 3 - (2nd stage) 

water to 
remove acid 

and dilute 

sodium hydroxide 

t/ ice-water 
r [.cooling 

t sodium hydrogen 
carbonate opt. 

Fig. 105. 
Preparation of sodium carbonate. 

Filter this off, wash the solid residue two or three times with a little 
cold water, and then transfer the solid to a dish and heat. Finally 
sodium carbonate will be obtained as a fine white powder. 

2NaHC0 3 -»- Na,C0 3 + H,0 + CO, 
Carbon dioxide is evolved during the reaction. 

Properties and uses of washing soda 

Washing soda is sodium carbonate decahydrate, Na,CO 3 .10H,O, 
large translucent crystals. 

Efflorescence. On exposure toairthecrystals lose weight and become 
coated with a fine white powder which renders them opaque. Each 
molecule of washing soda has given up to the atmosphere 9 mole- 
cules of water of crystallisation. 

Na.CO 3 .10H,O-»- Na a C0 3 .H,0 + 9H,0 
sodium carbonate sodium carbonate 
decahydrate monohydratc 

Such an action, that is, the giving up of water of crystallisation to the 
atmosphere, is termed efflorescence. 

The crystals are readily soluble in water. 

Anhydrous sodium carbonate can be made by heating the hydrated 
sodium carbonate. It is a fine white powder and does not dissolve as 
readily in water as do the crystals. 

Solutions of sodium carbonate in water are alkaline to litmus. 
This is due to the feebly acid properties of carbonic acid. This acid 
is expelled by almost every other acid, and hence, sodium carbonate 



acts like sodium hydroxide, although the former is a salt and the 
latter an alkali. 

Na a CO a + 2HC1 -»- 2NaCl + H a O + CO, 

Sodium carbonate can be used quantitatively in volumetric analysis, 
as if it were an alkali. 

Notice that sodium and potassium carbonates are both soluble in 
water and are not decomposed at a red heat. 

Sodium carbonate is used for the softening of water for domestic 
purposes, and in the manufacture of glass, borax, caustic soda and 
water-glass. It is a constituent of many 'dry soap' powders. 

Calcium carbonate, CaC0 3 

This occurs as limestone, marble, chalk, and in many other forms, 
and, since it is insoluble, can easily be made in the laboratory by 
double decomposition (see p. 233). It is seen as a white precipitate 
when carbon dioxide is bubbled into lime-water. 

Ca(OH), + CO a -► CaCOj + H a O 

The chalk can be obtained by filtering off the precipitate, washing 
it a few times with hot water and allowing it to dry. 

It is attacked by dilute hydrochloric and nitric acids with the 
evolution of carbon dioxide, for example: 

CaCOa + 2HC1 -*• CaCl, + H a O + CO, 

With dilute sulphuric acid, however, the action slows down and 
finally stops, particularly if the calcium carbonate is in lump form, 
for example, marble. The reason is that calcium sulphate, being only 
sparingly soluble, forms a protective layer on the outside preventing 
the sulphuric acid from acting upon the solid within. 

Although practically insoluble in pure water, calcium carbonate is 
dissolved by water which contains dissolved carbon dioxide, be- 
cause it forms soluble calcium bicarbonate (hydrogen carbonate). 

CaCO s + H a O + CO a -»- Ca(HC0 3 ), 


(See pp. 202-6 for treatment of calcium carbonate under 'hard- 
ness of water' and 'stalagmites and stalactites'.) 
Action of heat on calcium carbonate (see p. 244). 

Zinc carbonate, ZnC0 3 

This is formed as a white precipitate when sodium hydrogen 
carbonate solution is added to a solution of zinc sulphate in water. 


ZnS0 4 + 2NaHCO,-> ZnCO, + Na,S0 4 + H a O + CO a 
(If sodium carbonate is used, basic carbonate of zinc is formed.) 
The white precipitate is filtered off, washed with hot distilled water 
and allowed to dry. 

Zinc carbonate is attacked by dilute acids liberating carbon 
dioxide, for example: 

ZnCO, + 2HC1 ->- ZnCl, + H a O + CO, 
On heating a little zinc carbonate in a test-tube, carbon dioxide is 
given off and zinc oxide (yellow when hot, white when cold) remains 
in the test-tube. 

ZnC0 3 (c) ->• ZnO(c) + CO,(g) 
Zinc carbonate is used medicinally. 

Lead(II) carbonate, PbC0 3 

This is made in the laboratory, as a white precipitate, by adding 
sodium hydrogen carbonate solution to a solution of lead(II) nitrate 
in water. 

Pb(N0 3 ) a + 2NaHCO a -> PbCO a + 2NaNO s + H a O + CO, 
Sodium carbonate precipitates basic lead carbonate. 

White lead, basic lead carbonate, Pb(OH),.2PbC0 3 , is used ex- 
tensively as a paint when mixed with oils. It is made by subjecting 
strips of lead to the action of acetic acid, water-vapour, carbon diox- 
ide and air. It is poisonous and blackens rapidly in industrial areas, 
where hydrogen sulphide occurs. 

N.B. LeadfJI) carbonate is not readily acted upon by either dilute 
hydrochloric or sulphuric acids. A layer of insoluble chloride or 
sulphate formed round the carbonate protects it from further action. 
Dilute nitric acid attacks it to liberate carbon dioxide in accordance 
with the general action of acids on carbonates: 

C0 3 2 " + 2H+ 

H a O + CO, 

Copper(II) carbonate, CuC0 3 

This is made by double decomposition and it is usually obtained as 
a basic salt, having the formula CuC0 3 .Cu(OH) a . This is a bright 
green powder which liberates carbon dioxide on being heated, and 
black copper(II) oxide is left. 

CuCO a -»■ CuO + CO a ; Cu(OH), -*■ CuO + H a O 

It dissolves in dilute acids with the liberation of carbon dioxide, for 
example : 

CuCO, + H a S0 4 -*■ CuS0 4 + H a O + CO a . 



We are only concerned with the hydrogen carbonates of sodium 
and calcium, all the others being unstable or unimportant. 

Sodium hydrogen carbonate (baking soda), NaHCO . 

Sodium hydrogen carbonate is manufactured by saturating a wet 
mush of Solvay sodium carbonate and water with carbon dioxide. 
The product is washed with cold water and dried. 

CO a 8 - + H,0 + CO, -»■ 2HCO s - 

Laboratory preparation of sodium hydrogen carbonate 

Sodium hydrogen carbonate is made in the laboratory by bubbling 
carbon dioxide for some time through a concentrated solution in 
water of either sodium hydroxide or sodium carbonate (see Fig. 105, 
p. 313). In the latter case the reaction takes place much more quickly: 
CO, + Na,CO a + H,0 -► 2NaHCO s 

With sodium hydroxide 

2NaOH 4- CO,-*- Na,CO a + H,0 


then Na,CO a + H,0 + CO, ->■ 2NaHC0, 

In both cases the hydrogen carbonate is deposited as a white 
powder and this is filtered ofT, washed two or three times with a little 
cold distilled water, and allowed to dry. 

With dilute acids, the hydrogen carbonate liberates carbon 

HC0 3 - -f H + -*• H,0 + CO, 



Action of heat on sodium hydrogen carbonate 

Place a small amount of sodium 
hydrogen carbonate in a dry test- 
tube and heat gently with the lip of 
the tube projectinginto a boiling-tube 
containing lime-water (see Fig. 106). 
A gas is given off which turns the 
lime-water milky and is, therefore, 
carbon dioxide. Water is seen to 
condense on the cooler parts of the 

2HCO," -> C0 3 2 - + H,0 + CO, 

turns mi/ky 

Fio. 106. 

Action of heat on sodium 


The white residue is sodium carbonate. This test distin- 
guishes sodium hydrogen carbonate from sodium car- 
bonate, which is unaffected by heat. 

Sodium hydrogen carbonate is used in the manufacture of baking 
powders. Under the action of heat it decomposes, as above, and 
gives off carbon dioxide which causes the cake to 'rise' and so be 
light. This is why it is commonly called 'baking soda*. Baking powders 
also contain rice powder as a diluent, and tartaric acid (or a similar 
compound) to react with the sodium carbonate, which would other- 
wise be left when the hydrogen carbonate decomposes. 

Calcium hydrogen carbonate 

See Hardness of Water, p. 202. 

Its method of preparation and its reactions are chemically similar 
to those of sodium hydrogen carbonate. 

Calcium hydrogen carbonate cannot, however, be isolated as a 
solid, since it decomposes too easily, and all its reactions are carried 
out in solution. 


1. Graphite is called an 'allotropic form' of carbon. What do you under- 
stand by this statement? Give one other example of allotropy. How would 
you prove by a quantitative experiment that the statement is correct in the 
case of graphite and pure charcoal ? (N.U.J.B.) 

2. What is the action of (a) carbon dioxide; (6) steam, upon red-hot 
carbon and what is the practical importance of these reactions? (O. and 

3. What are the general methods employed for the preparation of 
metallic carbonates? 

5 g of a mixture of anhydrous sodium carbonate and sodium bicarbonate 
were heated until there was no further loss in weight. The resulting solid 
weighed 3.84 g. Find the percentage weight of the normal carbonate in the 
mixture. (Na = 23; C - 12.) (L) 

4. How does calcium carbonate occur in nature? Describe qualitative 
and quantitative experiments which you would do in order to determine 
the composition of this substance. (L.) 

5. Describe the preparation of a pure dry specimen of carbon monoxide. 
From carbon monoxide how would you prepare and collect a pure 

specimen of carbon dioxide? (C.) 

6. Describe how you would find by experiment the weight of carbon 
dioxide that can be obtained from 1 g of calcium carbonate. 

How would you prove that the gas obtained from calcium carbonate is 
identical with that produced when carbon is burnt? (O. and C.) 

7. Describe the preparation and collection of carbon monoxide and 
compare its properties with those of carbon dioxide. What would be the 
result, in the case of each oxide, of passing it (a) over heated copper(II) 



oxide; (b) oyer the red-hot charcoal ; (c) into cold, dilute sodium hydroxide 
solution? (B.) 

8. When oxalic acid is heated with concentrated sulphuric acid a mixture 
of the two oxides of carbon is evolved. 

Give a detailed account of how you would: 
ill tEfaS a s P ccimen of P urc carbon dioxide from this mixture. 
(b) Find the percentage by volume of each gas in the mixture. (O.) 

9. Describe carefully an apparatus by which you could collect for 
examination the products of the action of steam on red-hot carbon. 

Describe how you would identify each of the products of the reaction. 
Of what importance is the reaction in industry? (C.) 

10. How is carbon monoxide prepared ? Describe the chief properties of 
this compound. What volume will be found after the explosion of a 
mixture of one volume of carbon monoxide and ten volumes of air? (You 
may assume air to contain one-fifth by volume of oxygen.) (O. and C.) 

11. Describe how you would prepare sodium hydroxide from sodium 

State clearly how you could: 

(a) Distinguish sodium carbonate from sodium hydrogen carbonate; 

(p) Prepare sodium hydrogen carbonate from sodium carbonate. (O.) 

12. Write a short essay on •Flames', with special reference to those of 
hydrogen and coal-gas (luminous and non-luminous). (O.) 

13. Describe, using diagrams, the structure of the flame of a Bunsen 
burner. How do you account for the differences in the flame when the air 
holes of the Bunsen burner arc closed? Indicate experiments which might 
be carried out in support of your answer. (L.) 

14. What views were formerly held as to the nature of combustion, and 
what led to their overthrow? Who first gave the true explanation of com- 

State clearly all that you believe to be taking place in the burning of a 
candle. (L.) 

15. Starting with solid sodium hydroxide, describe how you would pre- 
pare in the dry solid state (a) sodium carbonate; (h) sodium hydrogen 
carbonate. Give two tests by means of which you could distinguish between 
these compounds. Describe and explain what happens when crystals of 
washing soda are exposed to the air. (L.) 

16. Explain fully any four of the following: (a) Why a Bunsen burner 
strikes back'; (b) the comparatively low percentage of carbon dioxide in 

the air; (t) the difference in the bleaching action of chlorine and sulphur 
dioxide; (rf) why sodium sulphate dissolved in water increases its con- 
ductivity ;(e) how the flameof a candlediffers from the flameof hydrogen. (L.) 

1 7. What is flame? By means of three labelled sketches only compare the 
structures of (a) a hydrogen flame; (b) a luminous Bunsen flame- (c) a 
candle flame. How would you demonstrate the presence of unburnt gas in 
a candle flame, and that this unburnt gascan be used asareducing agent ?(L.) 

18. Powdered marble is strongly heated in the open air until no further 
change occurs. Explain what happens and give three tests proving that the 
resultant residue is different from the powdered marble. Give short 
accounts of two important uses of the residue. Why does a precipitate 
form when temporarily hard water is boiled? How does the impurity enter 
the water in the first place? (L.) 

Chapter 24 

Organic Chemistry 

ORGANIC chemistry is the chemistry of the compounds of 
carbon; that is, all organic compounds contain carbon with, 
also, one or more other elements— hydrogen, oxygen, chlorine, 
nitrogen and metals. Carbon shows exceptional behaviour, in a 
chemical sense, by forming chains of its atoms (sometimes of very 
great length, e.g., about 2000 carbon atoms in polythene) and rings 
of its atoms. Sometimes, the chains and rings are included in the same 
molecule. Consequently, carbon produces a very great number of 
compounds, many of them very complex, and it has become neces- 
sary to make a separate study of these compounds as organic chem- 
istry. Nominally, every compound of carbon is an organic compound. 
For historical and conventional reasons, however, a few of the simpler 
carbon compounds, such as carbon dioxide and sodium carbonate, 
are usually studied with non-carbon compounds in inorganic chem- 
istry. Examples of chain and ring formation by carbon atoms will be 
found in the present chapter. 


Hydrocarbons are compounds containing hydrogen and carbon and 
no other element. That is, a hydrocarbon has the molecular formula, 
C x H y , x and y being whole numbers. For example, methane, CH 4 , 
ethylene, C S H 4 , and benzene, QH 4 , are hydrocarbons. The last is the 
liquid known to the motor trade as benzol. 

Paraffin Hydrocarbons or Alkanes 

The members of this group of hydrocarbons (also called simply 
paraffins) are distinguished by possessing the general molecular 
formula, C,,H M+8 , where n is 1, 2, 3, etc., for successive members of 
the group. This general molecular formula will be justified after the 
properties of a typical member of the series have been studied. The 




first member of the series {n = 1) is methane, CH 4 , and the second 
(n = 2) is ethane, C,H 4 . Both are gases at room temperature and 

Laboratory preparation of methane 

Sodium acetate (anhydrous) is ground with its own weight of 
soda-lime. The mixture is heated in a hard-glass flask (Fig. 107). 
Methane is evolved and collected over water. When heating is 
finished, the delivery tube must be removed from the water at once; 
otherwise, water may suck back into the hot flask and cause an 
explosion. Soda-lime acts as a source of sodium hydroxide. 
CH 3 COONa(c) f NaOH(c) -> Na 2 C0 3 (c) + CH 4 (g) 

Sodium Acetate and 
Soda - time 


Fig. 107. 
Preparation of methane. 

Properties and reactions of methane 

Physical properties 

Methane is a colourless gas with no smell. It is almost insoluble 
in water and much less dense than air in the same conditions (V.D. of 
methane is 8, of air is 14.4). 

Chemical reactions 

(I) Combustion. Methane burns (or explodes) in air on the appli- 
cation of a flame (or electric spark). It produces carbon dioxide and 

CH 4 + 20 2 -y C0 2 + 2H 2 0; A// = - 79 kJ g-eqn" 1 
This is an exothermic reaction, utilised as a means of industrial and 
domestic heating when methane is burnt as a constituent of coal-gas. 
Similarly, other (liquid) alkanes are used as petrol or fuel oil in 
exothermic combustion with air. For example, pentane, a constituent 
of petrol, burns in the following way. 

QH I2 + 80 2 -> 5CO t + 6H 2 with heat evolved 


Methane is the important constituent of natural gas which usually 
accompanies petroleum deposits. This gas is at present being im- 
ported into Britain in pressurised tanker ships, as liquid, from the 
Saharan oilfields and added to coal-gas supplies. Recent borings in 
the North Sea have revealed reserves of natural gas estimated (May 
1969) to yield at least 3000 million cubic feet of the gas per day for 
many years; this daily output would meet the whole of Britain's 
current need of gaseous fuel. Several years will be required for laying 
pipe-fines and for modifying or replacing gas appliances to corres- 
pond with the fact that, volume for volume, natural gas has twice 
the calorific value of present gas supplies. Natural gas is non-poison- 
ous and has no smell, unless hydrogen sulphide is present. 

(2) With chlorine. Methane reacts (slowly at ordinary temperature) 
with chlorine, the reaction being catalysed by light (photocatalysis). 
The first product is monochloromethane, CH 3 C1 (or methyl chloride, 
the group, CH 3 _ , being known as methyl). 

CH 4 + Cl 2 ->- CH 3 C1 + HC1 

In a similar way, excess of chlorine may produce, with increasing 
difficulty, dichloromethane, CH 2 C1 2 , chloroform, CHC1 3 , and carbon 
tetrachloride, CC1 4 . 

These are called substitution reactions because chlorine successively 
replaces hydrogen in the methane molecule. Notice that the hydrogen 
is expelled in combination with chlorine as HCI, not as free hydrogen. 

The above two reactions are the only ones given by methane or 
other paraffins in ordinary conditions, i.e., without the use of high 
temperature and pressure or other exceptional means. That is, the 
paraffins are a rather inert group; in fact, their name was coined 
from the Latin, parum, little, and affinis, affinity. 

Structure of alkanes 

In the alkane, all carbon atoms exercise a covalency of 4. The 
simplest alkane has, therefore, the molecular formula, CH 4 , and 
is methane. Each covalent bond represents a shared pair of electrons, 
one each from the carbon atom and a hydrogen atom. In more 
complex alkanes, carbon atoms form chains by combining together 
by covalency. With chains of two carbon atoms {ethane, C 2 H g ) and 
three carbon atoms {propane, C 8 H 8 ), the molecular structures are 
shown below, together with methane. 

H H H 

I 1 1 

H— C— H H— C— C— H 

k kk 

methane ethane propane 

H H H 
H— C— C— C— H 






Much longer chains are also produced, such as the following: 


In such structures, it will be seen that all the carbon atoms except 
the two at the ends of the chain use two units of valency to attach 
themselves to the carbon atoms on each side, so forming the carbon 
chain. Each carbon atom (except the end ones) then has two units of 
valency left to combine with two hydrogen atoms. That is, except 
for the end carbon atoms, the relation of carbon to hydrogen is 
C n H sn . But each of the end carbon atoms uses only one valency unit 
in forming the chain and so can combine with an extra hydrogen 
atom, so producing the general molecular formula, C„H 2n+2 . Many 
alkancs have a branched carbon chain such as the following. 

H H H H H H H 


H H H H H H 

H— C— H 

H— C— H 


In this case, a hydrogen atom is lost from the carbon atom on which 
the branching occurs. But it is restored at the open end of the branch, 
so the general molecular formula is unchanged. 

American petroleum consists largely of these alkanes from pen- 
tane C 5 H,., to the very complex compound, C 43 Hgg. The four 
simplest alkanes are all gases in ordinary conditions, i.e., CH 4 -C 4 H 10 . 
Butane, C 4 H 10 , is supplied in pressurised containers as color gas. 

Homologous series 

A series of compounds related to each other as the alkanes are 
(above) is called a homologous series. Such a series has the following 

(1) All members conform to a general molecular formula, e.g., 
for alkanes, C„H W _ S . 

(2) Each member differs, in molecular formula, from the next by 
CH„ e.g., alkanes are CH 4 , C 2 H„, C 3 H 8 and so on. 

(3) All members show similar chemical reactions, though varying 

in vigour. For example, all alkanes burn in air and give substitution 
reactions with chlorine (p. 320). 

(4) The physical properties of members change gradually in the 
same direction along the series, e.g., in the alkanes, boiling-points 
and freezing-points rise (CH 4 — a gas; C s H ia - a liquid; C8oH 4a — a 
solid, at ordinary temperature and pressure). Also, densities increase 
and solubility in water decreases as the number of carbon atoms per 
molecule increases. 

(5) General methods of preparation are known which can be 
applied to any member of the series. 

Other homologous series are alkene hydrocarbons, C„H 2 „, alcohols, 
C„H 2B+1 OH, and fatty acids, C„H 2n+1 CO a H. These series are con- 
sidered in the present chapter. 

Cyclic compounds 

In addition to producing open chains of atoms, as in the alkanes, 
carbon can produce closed rings of atoms. Compounds possessing 
such carbon rings arc said to be cyclic 
compounds. The commonest rings are 
those composed of five or six carbon 
atoms. This is so because the four 
valencies of the carbon atom are direc- 
tional and distributed symmetrically in 
three-dimensional space (Fig. 108). This 
gives a natural angle of about 109° 
between any pair of valency directions. 
The angles in a regular five-membered 
ring are 108° and in a regular six- 
membered ring are 120°. That is, the 

natural angle between the carbon valency directions needs very little 
strain to form 5- or 6-membered rings. 

The most important is the 6-membered ring of benzene, C„H 6 , and 
its derivatives. This is usually represented with alternate single and 
double carbon-to-carbon valency bonds (Kekule's formula), though 
this representation is known to be unsatisfactory in some respects. 

.about 109° 

Fio. 108. 

Valency directions of the 
carbon atom. 

Derivatives of benzene are very important chemicals, e.g., aniline, 
C g H 6 NH B , and phenol, C,H 6 OH, both of which can be obtained 



from coal-tar as source. More complex benzene derivatives include 
the M. and B. drugs, aspirin, and a wide range of dye-stuffs. Cyclic 
compounds may also contain straight carbon chains, too, which are 
known as side-chains in this case. An example is the following. 


HC/\C-CH,-CH 8 -CH 1 -CH 8 NH 1 



It is obvious that such combinations may become very complex. This 
is why carbon appears to be the only element which can supply the 
very varied and complex compounds needed by living organisms to 
carry on their vital processes. 


American petroleum consists mainly of alkanes from QH 12 , pen- 
tane, to C^Hgg, usually associated with salt water and the gaseous 
paraffins. After removal of impurities (mainly sulphur compounds), 
the petroleum is distilled into four fractions: 

Up to 200°C Naphtha 1 QH lt -C lt H M 

20O-260°C Kerosene C 12 H M -C 1S H 31 

26O-300°C Gas-oil C 16 H 3! -C 17 Hs, 

Above 300°C Fuel oil 

Redistillation of naphtha yields petrol, b.p. 50°-60°C, mainly C 5 H 1S 
to C,H g0 . Other products of redistillation of the other fractions are 
paraffin, lubricating oil, petroleum jelly and paraffin wax. Much of the 
gas-oil fraction is cracked to yield more petrol, i.e., the complex 
molecules are broken down into simpler units. At about 12 atm pres- 
sure and 520°C, gas-oil yields petrol hydrocarbons and valuable 
alkene by-products, e.g., ethylene, C,H«, and propylene, C 3 H 6 . (For 
fractional distillation, see p. 218). 

C ls H at — * CgHjg -f 2C a H« -f C 3 H e 
Petroleum refining occurs usually on an estuary, e.g., at Fawley 
(Southampton), to avoid land transport of bulky imported oil and 
to secure large flat areas for its storage, segregated because of fire 


Isomerism occurs often among organic compounds because of 
their complexity. 

1 For naphtha as a source of fuel gas, see p. 292. 


Isomerism is the occurrence of two or more compounds 
with the same molecular formula but different molecular 

Isomers of the same molecular formula have different physical and 
chemical properties because of structural differences. 

Isomers of molecular formula C 4 H 10 
H H H H 

I J 

H H H H 

orCH 3 .CH 2 .CH 2 .CH 3 

normal butane; a straight carbon 

H— C 



H H H 
— A-C-H 



or CH 3 .CH(CH 3 ).CH 3 

isobutane; a branched carbon 

These compounds differ in boiling-point, freezing-point and density. 
Their chemical reactions are not very different because they both 
belong to the alkane homologous series. 

Isomers of molecular formida C 2 H O 
H H 
H-C— C-O-H 


H H 

H-C-O-C— H 



or CH 3 .CH 2 .OH 

ethanol; an alcohol. The oxygen 
atom is part of a hydroxyl 
group, OH. 

or CH 3 .O.CH 3 

dimethyl ether; the oxygen 
atom forms a bridge between 
the two carbon atoms. There 
is no hydroxyl group. 

Ethanol is a liquid, dimethyl ether a vapour at room temperature and 
pressure. As an alcohol, ethanol reacts rapidly with sodium and 
phosphorus pentachloride (p. 335). The ether has neither of these 


Isomers of molecular formula C 2 H 4 CI 2 
H H 

CI— c— c— ci 


or CH 2 C1.CH 2 C1. 

One chlorine atom is combined 
with each carbon atom. 

H H 
H— C— C— CI 


or CH 3 .CHCI 2 . 

Both chlorine atoms are com- 
bined with the same carbon 

These compounds differ in boiling-point and freezing-point. If 
hydrolysed, the left-hand one yields glycol, C 2 H 4 (OH) 2 ; the other 
yields aldehyde, CH 8 .CHO. 

define Hydrocarbons or Alkenes 

The define hydrocarbons (or alkenes) are members of a homo- 
logous series of general molecular formula, C„H 2B . That is, each 
member of the series has two fewer hydrogen atoms per molecule 
than the corresponding alkane, e.g., alkane, C 8 H 8 , alkene, C s H a . The 
reason for this will be mentioned later. The most important alkene 
is ethylene, C 2 H 4 , which is a gas at room temperature and pressure. 

Laboratory preparation of ethylene (or ethene) 

Ethylene is usually prepared by the dehydration of ethanol (alcohol) 
by hot, concentrated sulphuric acid. To ethanol (50 cm 3 ), slowly with 


Cong. Sulphuric 
acid and ^J 


Fio. 109. 

Preparation of ethylene. 


shaking and cooling under the tap, is added concentrated sulphuric 
acid (100 cm 3 ). Apparatus is set up as in Fig. 109, and the mixture is 
heated with care to about 180°C. Ethylene is evolved and collected 
over water. The wash-bottle of alkali solution serves to remove sul- 
phur dioxide, a by-product produced in small amount as the alcohol 
reduces the sulphuric acid slightly. 
The alcohol and acid first produce ethyl hydrogen sulphate. 
C.H.OH + H 2 S0 4 -> C 2 H 5 .HS0 4 r- H a O 
When heated, ethyl hydrogen sulphate releases ethylene. 

C a H s .HS0 4 -> H 2 S0 4 + C 2 H 4 
These two reactions are equivalent to the dehydration of alcohol by 

the acid. 

C 2 H s OH-H 2 0->-C 2 H 4 

Industrial production of ethylene 

In industry, ethylene is produced by cracking the gas-oil fraction 
from distillation of petroleum. At 12 atm pressure and 520°C, alkanes 
in the petrol range (C S H, S C B H 20 ) are produced, together with 
ethylene and, possibly, other alkenes. For example, 
C„H M -* C 8 H 18 + 3C 2 H 4 + C 8 H 6 . 

Properties and reactions of ethylene 

Physical properties. Ethylene is gaseous at room temperature and 
pressure, colourless, almost insoluble in water and slightly less dense 
(V.D. - 14) than air (V.D. = 14.4). 

Chemical reactions 

(1) Combustion. Ethylene burns (or explodes) in air if a light (or 
electric spark) is applied. The products of complete combustion are 
carbon dioxide and steam, but the flame tends to be smoky from 
unburnt carbon because of its high proportion (about 86%) in ethy- 

C 2 H 4 + 30. -»- 2CO, + 2H t O 

(2) Addition reactions. Ethylene gives a number of addition re- 
actions in which two hydrogen atoms (or their equivalent) are taken 
into combination per molecule of ethylene to form a single product. 
Ethylene is, therefore, said to be unsaturated. Structurally, this means 
that it contains a double bond between carbon atoms, H 2 C ^CH 2 . 
This will be further considered later. 

(a) With chlorine or bromine. With chlorine at ordinary tempera- 
ture, ethylene combines rapidly to give an oily liquid, ethylene 


H 2 C=CH 2 + CI,-* CH.Cl.CH.Cl 





This accounts for the name, defiant gas, by which ethylene was once 
known, from Latin, oleum, oil, andy?o, I become. The term, define, 
has developed from the older name. 

The similar reaction with bromine vapour rapidly destroys the 
reddish-brown colour of the vapour. This acts as a distinguishing 
test between ethylene and gaseous paraffins, e.g., methane or ethane, 
which do not give this rapid colour change. 

H 2 C=CH, + Br 2 -> CH 2 Br.CH 2 Br 

(b) With hydrogen iodide. Ethylene combines rapidly with hydro- 
gen iodide (vapour) at ordinary temperature to produce ethyl iodide. 

H,C=CH 2 + HI -J- CH 3 .CHJ(or C 2 H 6 I) 

The group, C,H 6 -, is called the ethyl group. The gases, HBr and HC1, 
combine similarly but more slowly. 

(c) With concentrated sulphuric acid. Ethylene is absorbed rapidly 
by this acid at room temperature to form ethyl hydrogen sulphate. 

H 2 C=CH, + H 2 S0 4 -»- CH 3 .CH 2 HS0 4 (or QHb.HSOJ 

The action is reversed at about 180°C, liberating ethylene. 

(d) With hydrogen. Ethylene combines with hydrogen if the two 
are passed over finely divided nickel (catalyst) at about 200°C. The 
product is the alkane, ethane. 

H 2 C=CH, + H 2 -*- CH 3 .CH S (or C 2 H,) 

This (Sabatier's) reaction is important as representing the essential 
change in the conversion of oils into margarine. An oil possesses a 
long carbon chain (15-17 atoms) in which there is one ethylenic 
double bond,— CH=CH— . Pressure is used (5 atm) and a somewhat 
lower temperature (180°C) and the oil is hydrogenated in the presence 
of nickel. This saturates the double bond, so converting the oil into 
a fat, which is then sold as margarine. 

— CH=CH— + H 2 -* — CH a — CH»— 

(e) With potassium permanganate. Ethylene reacts rapidly if shaken 
with potassium permanganate solution at ordinary temperature. If 
acidic, the solution is decolorised, being reduced to a manganese salt; 
if alkaline, it turns green, being reduced to a manganate. In both cases, 
the ethylene is converted to ethylene glycol by water and oxygen from 
the permanganate. 

H,C=CH, + H 2 + (O) ->• C,H 4 (OH), 

This reaction distinguishes ethylene from all gaseous alkanes, e.g., 
methane or ethane, which have no reaction with permanganate. 
Ethylene glycol is the material used in anti-freeze solutions for motor- 
car radiators and for the production of terylene (p. 337). 



If ethylene, with a trace of oxygen present, is pressurised to about 
1000 atm and heated to start the action, it is polymerised to form 
polyethylene or polythene. Polymerisation is the combination of two 
or more molecules of a given compound to form one complex mole- 
cule with no gain or loss of material. 

nA— *■ A„ 

The product, A„, is called a polymer of the original compound, A. 
In the case of ethylene, the polymerisation is quite exothermic and, 
once the reaction has started, cooling is required. 

Polythene is very resistant to the common types of chemical action 
and can be moulded (while hot) into a great variety of domestic and 
scientific articles- buckets, bowls, bags, flexible containers, funnels; 
wash-bottles. The polymerisation can be stated as: 

3n(CH 2 =CH 2 ) ->- (-CH 2 -CH 3 -CH 2 -CH 2 -CH 2 -CH 2 -)„ 

where n is about 300. From this, polythene resembles a highly com- 
plex alkane and shares the chemical inertness of this group. 

Unsaturation in ethylene 

Ethylene is said to be unsaturated because one molecule of it can 
combine with two hydrogen atoms (or their equivalent) in addition 
reactions. In general, 

H 8 C=CH 2 + A.B -► AH 2 C-CH 2 B 

where A and B are univalent atoms or groups. Several examples of 
these reactions were given in the recent sections of this chapter. The 
C to C double bond which represents this unsaturation indicates that, 
between the carbon atoms concerned, two pairs of electrons are 
shared in covalency. 

H H 


H H 

oo o oo 

corresponds to CSC 

oo o oo 

H H 

When the addition reaction has taken place, e.g., with Cl 2 , only two 
electrons are shared between the carbon atoms; only single covalent 
bonds are present in the molecule, the compound is said to be satur- 
ated and is an alkane derivative. It gives no addition reactions but 
reacts by substitution. The addition of chlorine to ethylene is shown 
by the following equation. 



H H 



H H 


+ CI a 

CI— C— C— CI 


Acetylene or Ethyne 

Acetylene is really the first member of a homologous series of 
hydrocarbons called the alkynes. The general molecular formula of 
the series is CH^-,. For acetylene itself, n = 2, and the molecular 
formula is C 2 H 2 . So far, acetylene is the only member of the series 
which has acquired general importance. Alkynes are characterised 
by possessing a carbon-to-carbon triple bond at one point in the 
carbon chain. The triple bond involves the loss of four hydrogen 
atoms per molecule in an alkyne hydrocarbon when compared with 
the corresponding alkane. 

H— Ci=C— H 


H H 


The triple bond involves the sharing of three pairs of electrons in 
covalency between the two carbon atoms. 

Preparation of acetylene 

For apparatus, see Fig. 110. Cold water is dripped on to calcium 
carbide. Much heat is evolved and sand protects the flask from break- 




Copper Sulphate 
in dil. Sulphuric 



Fio. 110. 
Preparation of acetylene. 


age. The chief impurity phosphine, PH 3 , is absorbed by the acidified 
copper(II) sulphate solution. Acetylene is collected over water. 
CaC 2 (c) + 2H 2 -> Ca(OH) 2 + C a H 8 (g) 

Properties of acetylene 

Physical. In ordinary conditions, acetylene is a colourless gas, 
almost insoluble in water and having a sweet smell when pure. It is 
slightly less dense than air (V.D. is 13; V.D. of air is 14.4). Acetylene 
is strongly endothermic. 

2C + H, -> QH t ; AH = + 200 kJ g-eqn" 1 
Like endothermic compounds in general, it is rather unstable and 
liable to explode if stored under compression alone. For industrial 
purposes, it is stored in steel cylinders in solution in acetone at about 

12 atm pressure. 

Chemical reactions. Having a carbon-to-carbon triple bond in its 
structure, acetylene is an unsaturated compound. It gives a number of 
addition reactions, combining with four hydrogen atoms per molecule 
(or their equivalent) as a maximum, then producing the saturated 
alkane, ethane, or one of its derivatives. These reactions are of the 


+ 2AB —*■ | where A and B are univalent 


Notice that the two A groups usually combine with the same carbon 
atom, as do the two B groups. Compounds of the type CHAB. 
CHAB are rarely produced. 

(a) With bromine. At ordinary temperature, acetylene combines 
rapidly with bromine, forming acetylene tetrabromide. 

CH CHBr 2 

HI +2Br t ->| 

CH CHBr 2 

Chlorine gives a corresponding reaction provided it is diluted with 

an inert gas, but, if pure acetylene and chlorine are mixed, a violent 

explosion occurs, with formation of carbon and hydrogen chloride. 

C 2 H f + Cl 2 -»-2C + 2HCl 

(b) With hydrogen. If passed with twice its own volume of hydrogen 
over nickel (catalyst) at about 200°C, acetylene forms ethane. 

C,H 2 + 2H 2 -*C 2 H 8 

(c) With halogen acids (as gases). Acetylene combines readily with 
hydrogen iodide at room temperature to form ethylidene iodide. 


+ 2HI ->- 1 

:H ch, 




A corresponding reaction is given by hydrogen bromide at I0O°C, 
but the reaction with hydrogen chloride is very slow. 

(d) With water. Acetylene reacts additively with water if passed 
into dilute sulphuric acid at about 96°C with mercury(II) sulphate 
present as catalyst. The product is aldehyde (acetaldehyde). 

C 2 H 2 + H 2 -y CH3.CHO 

This is the first stage of the manufacture of acetic acid from acetylene 
(p. 338). 

(e) With acidified potassium permanganate solution. At room tem- 
perature, with shaking, acetylene quickly decolorises this solution 
(i.e., reduces it) with formation of oxalic acid. 

8Mn0 4 ~ + 5C a H a + 24H+ -*• 5H 2 C 2 4 + 8Mn 2+ + 12H 2 
Notice that acetylene decolorises bromine vapour (see (a) above) 
and decolorises acidified permanganate solution (see (e) above). These 
reactions distinguish it from alkanes such as methane and ethane 
which do not give these changes; they do not distinguish acetylene 
from ethylene which also produces the same colour changes (p. 327). 

Metallic derivatives of acetylene 

Acetylene produces several metallic derivatives, cuprous or copper(I) 
acetylide and silver acetylide being the most important. 

Copper 1 1) acetylide. If acetylene is passed at room temperature into 
a solution of cuprous chloride in ammonia, a reddish-brown precipi- 
tate of copper(I) acetylide is at once thrown down. 

C,H, + 2CuCl + 2NH 4 OH -»• Cu 2 C 2 + 2NH 4 CI + 2H 2 
Silver acetylide. If acetylene is passed at room temperature into a 
solution of silver oxide in ammonia, a yellow precipitate is produced, 
rapidly turning grey. It is silver acetylide. 

C 2 H„ + Ag s O -»• Ag 2 C 2 + H,0 
If allowed to dry and heated, both acetylides are explosive, the silver 
salt much more violently. Both give off acetylene if warmed with 
dilute acid. 

Ag 2 C 2 + H 2 S0 4 ->- Ag 2 S0 4 -f C t H, 

The two reactions above distinguish acetylene from ethylene which 
forms no metallic derivatives. 

Polymerisation. Acetylene polymerises into the cyclic hydrocarbon, 
benzene, if passed through a tube at red heat. 

3C 2 H 2 — > C,H„ 

Combustion. Acetylene burns in air if ignited by a flame or electric 
spark and mixtures of the two can explode very violently. The flame 
is usually very sooty; acetylene contains over 90% of carbon and 



much of it remains unburnt unless special means are used to provide 
a good air supply. 

2C 2 H 2 + 50 t -> 4CO a + 2H s O 
A flame of acetylene burning in oxygen (the oxyacetylene flame) is 
very hot and can give a temperature of 2200°C. It is used in welding 
and in cutting up steel scrap. 

Manufacturing uses of acetylene 

Acetylene has recently become very important as a source of organic 
chemicals on the large scale, though some of the products made from 
it are very complex. The manufacture of aldehyde and acetic acid from 
acetylene is given on p. 338. It is also polymerised to ethyl acetate, 
using aluminium ethoxide as catalyst at 0°C. 

2CH3.CHO -> CH 3 .COOC s H 8 
Acetylene can also be converted to acetic anhydride, which, together 
with acetic acid (also an acetylene product), converts cellulose to 
cellulose acetate which is used as celanese rayon and cellophane. 
Acetylene is also the starting material for making polyvinyl chloride, 
which is used in electrical insulation and waterproofing, and polyvinyl 
acetate, which is used in sheets as the middle layer of Triplex safety 
glass. The uses of the oxyacetylene flame have been mentioned above. 


The alcohols form a homologous series of general molecular for- 
mula C„H 2n+1 OH (or C n H ani2 0). The first of these formula; is usually 
preferred because it shows the presence of the hydroxyl group, OH, 
which is the characteristic group of the alcohols. The first two mem- 
bers of the series are methanol, CH 3 OH, also known as wood spirit 
because of its early production by distillation of wood, and ethanol, 
C 2 H s OH, the alcohol of ordinary life. The two are also known 
(respectively) as methyl alcohol and ethyl alcohol. 

Manufacture of ethanol 

(1) By fermentation. The source of the alcohol is some form of 
starch; in western Europe, potatoes are commonly used. They are 
pressure-cooked to release the starch granules and then treated with 
malt (partially sprouted barley) for an hour at 60°C. Malt supplies an 
enzyme, i.e., an organic catalyst, called diastase. In these conditions, 
the starch is hydrolysed to a sugar, maltose. The empirical formula 
of starch is C 8 Hi O 5 . 

2C,H 10 O s + H 2 -»- C 12 H 22 O n (maltose) 


At room temperature, yeast is added. One of its enzymes, maltose, 
catalyses the hydrolysis of maltose to glucose. 

C„H u O u + H.O ->■ 2C„H I2 8 

Another enzyme of yeast, zymase, catalyses the decomposition of 
glucose to ethanol and carbon dioxide. 

C 6 H 12 0, -*- 2C 2 H s OH + 2CO, 

The 'wash*, containing less than 11% of ethanol, is converted by dis- 
tillation in a very efficient still to a liquid containing about 95% of 
ethanol and this is the purest form (surgical spirit) in which the 
alcohol is usually sold. The material known as methylated spirit con- 
tains about 85% of ethanol (with water) and has paraffin and colour- 
ing matter added to prevent its being drunk or used for chemical 
manufacture to evade excise duty. 

Fermentation of starch (or the sugar of grapes) produces many 
alcoholic beverages, e.g., varieties of beer and wines with 8-20% of 
alcohol; fermented and distilled products are known as spirits— 
whisky, gin, brandy, rum— and contain about 30% of alcohol. 

The effect of alcohol on human capabilities is, of course, attracting 
much contemporary attention, especially in its relation to motoring. 
The fact appears to be that even a little alcohol tends to impair skill 
and cloud judgement, though the extent of these effects depends on 
the individual and his alcoholic experience. Greater quantities of 
alcohol lead to much greater lack of muscular control, e.g., the 
drunken stagger, and, ultimately, to coma, the state of dead-drunk. 
A sufficient large alcoholic consumption may lead to continuance of 
the coma to the point of death. Methanol is more poisonous still: 
even moderate amounts may affect the optic nerve and produce blind- 
ness for several days, while greater amounts may produce permanent 
blindness or prove fatal. 

The breaking down of a complex material such as starch into much 
simpler substances, e.g., by fermentation into alcohol and carbon 
dioxide, is known as degradation and is a common way of utilising 
natural products. The distillation of coal in air-free retorts to yield 
coal-gas, tar, coke and ammonia liquor is another example of de- 

(2) Manufacture of alcohol from petroleum products. The cracking 
of petroleum oils to produce ethylene was described on p. 327 If 
ethylene is absorbed at 80°C under slight pressure in concentrated 
sulphuric acid, ethyl hydrogen sulphate is formed. If this product is 
diluted with water and distilled, ethanol is obtained. 



C,H« + H g S0 4 
C a H 5 .HS0 4 + H.O 

QH 5 .HS0 4 
CjHjOH + H a S0 4 

In a newer process, ethylene and steam are passed over phosphoric 
acid (catalyst) at 600°C and 60 atm pressure. The greater part of the 
world's industrial ethanol is now produced by these two processes. 
The alcohol is used as a solvent in stains and polishes and in the 
manufacture of essences, perfumes and drugs. Its use in beverages and 
methylated spirit was mentioned above. 

Properties of alcohols 

Physical. The simpler alcohols, e.g., methanol and ethanol, are 
liquids at room temperature and pressure, volatile, colourless and 
possessed of a characteristic smell. They mix with water in all pro- 
portions. Some alcohols are known with a carbon chain of about 
20 atoms. These are solids resembling paraffin wax in appearance. 

Chemical reactions. (1) Reactions of the hydroxy I group. All alco- 
hols contain the hydroxyl group, OH, and show corresponding 

(a) With sodium. If a small piece of sodium is added to 1-2 cm 3 of 
ethanol at room temperature, effervescence will occur with liberation 
of hydrogen (explosion on the application of a flame). Sodium ethoxide 
will eventually remain as a white, deliquescent solid. Methanol 
behaves similarly. 

2C 8 H 6 OH + 2Na -»- H a + 2C a H 6 ONa (sodium ethoxide) 
2CH a OH + 2Na -*■ H 8 + 2CH 3 ONa (sodium methoxide) 

(b) With a chloride of phosphorus. If a little phosphorus penta- 
chloride is added to 1-2 cm 3 of ethanol at room temperature, a 
vigorous reaction occurs with liberation of 'steamy' fumes of hydrogen 
chloride. The organic product is ethyl chloride, which escapes as 
vapour. This is the recognised test for the hydroxyl group, OH, in a 
straight-chain organic compound. Its effect is to replace the OH 
group by CI. 

C a H s OH + PC1 B ->- C 2 H 5 C1 + POCl 3 + HC1 

Methanol, CH s OH, gives a similar reaction to produce methyl chlor- 
ide, CH 3 C1, which is also a vapour. 

Phosphorus trichloride gives a similar reaction, but usually less 

3C 2 H 5 OH -f 2PCl 3 -> 3C S H 6 C1 + P,0 3 + 3HC1 

All members of the alcohol homologous series give the reactions (a) 
and (b) above, though with decreasing vigour as the number of 
carbon atoms increases. 

(2) Reaction with sulphuric acid. Ethanol reacts with sulphuric acid 
in two distinct ways. If the acid is in considerable excess and the 





temperature is about 180°C, the product is the gaseous alkenc, ethy- 
lene (p. 326). 

C a H s OH + H 2 S0 4 -> C 2 H 6 .HS0 4 + H s O 
C 2 H 6 .HS0 4 -»- C 2 H 4 + H 8 S0 4 

With a smaller proportion of the concentrated acid and a lower 
temperature (145°C), the chief product is ether (diethyl ether). This 
is a very volatile, sweet-smelling liquid, formerly much used as an 

C 2 H & OH + H 2 S0 4 ->• C 2 H 5 .HS0 4 + H 2 

Then, with excess alcohol available, 

C„H B .HS0 4 + HOC 2 H 5 -»- C 2 H 6 .O.C 2 H 8 + H 2 S0 4 . 

These are both dehydrating actions of sulphuric acid. The severer 
conditions (excess of the acid and a higher temperature) extract H 2 
from one molecule of the alcohol. 

C 2 H s OH - H 2 -»- C 2 H 4 

Milder conditions (relatively less acid and a lower temperature) 
extract H.O between two molecules of the alcohol. 

2C 2 H 6 OH - H 2 0-»- C 2 H 6 .O.C 2 H 6 

(3) Oxidation. Ethanol can be oxidised in two stages. If 1-2 cm 3 of 
ethanol are heated with potassium dichromate solution and dilute 
sulphuric acid, the liquid turns green (by reduction of the dichromate) 
and aldehyde is given oiTas a very acrid vapour. 

CH 3 .CH 2 OH -> CHj.CHO + 2H* + 2e~ 

Aldehyde can be further oxidised by more heated potassium dichrom- 
ate in dilute sulphuric acid and the product is acetic acid, the acid of 

CH 3 .CHO + H 2 0-»- CH3.COOH + 2H+ + 2e~ 
No further oxidation is possible except by combustion of acetic acid 
vapour to carbon dioxide and water. Because of its oxidation product 
(acetic acid), aldehyde is also known as acetaldehyde. 

Both these stages of oxidation can also be brought about by passing 
the alcohol or aldehyde, as vapour, with air over a heated catalyst. 
For the first stage, a copper coil is a suitable catalyst. 

CH 3 .CH 2 OH + iO s -»- CH 3 .CHO + H t O 
For the second stage, manganese(II) acetate is catalytic. 
CH 3 .CHO + ±0 2 -> CH 3 .COOH 

(4) Ester formation. Into a dry test-tube, put about 2 cm 3 of glacial 
{i.e., concentrated) acetic acid. Add a few drops of concentrated sul- 
phuric acid (catalyst) and shake. Add 4 cm 3 of ethanol, shake and 

place the test-tube in a beaker of cold water. Heat it slowly and keep 
the temperature as high as is possible without allowing the mixture 
in the test-tube to boil. After 20-25 minutes, pour the contents of 
the test-tube into an evaporating dish about half-full of a concen- 
trated common salt solution. Oily drops of liquid should rise to the 
surface (the ester, ethyl acetate), having a pleasant smell of pear- 

CH 3 .COOH + HOC 2 H 5 ^ CH 3 .COOC,H 6 + H a O 

Notice that the ester is formed from the acid by replacing ionisable 
hydrogen (of the COOH group) by the organic ethyl group from the 
ethanol. All the common alcohols produce esters with organic acids 
in this way; the esters usually have a pleasant fruity smell and find 
uses as flavouring materials and in perfumes. The production of a 
more complex product like the ester from simpler materials like the 
acid and alcohol is called synthesis and is a very frequent organic 

It can be reversed, to recover the acid and alcohol again, by 
hydrolysis, theoretically by water but, in practice, by boiling with 
dilute mineral acid (HC1 or HgSOJ or with aqueous caustic alkali as 
catalysts. In the acidic case, the products are the alcohol and acid, 
from which the ester is derived. 

CH 3 .COOC s H 6 + H s O ^ CH 3 .COOH + C 3 H 6 OH 
In the alkaline case, the above reaction occurs and is then followed by 
the production of the sodium (or potassium) salt of the acid, in this 
instance, sodium acetate. 

CH 3 .COOH + NaOH -»• CH 3 .COONa + H 2 

The production of soap by boiling a fat or oil with aqueous caustic 
soda is alkaline hydrolysis of an ester. The ester is the fat or oil, the 
alcohol produced is glycerol and the sodium salt is the soap. 
Fat or oil + caustic soda — >■ soap -f- glycerol 

The chemical formulae are very complex. 

Synthesis of terylene. Terylene is a very complex ester (a polyester) 
synthesised from the alcohol, glycol, C 2 H 4 (OH) 2 , and lerephtltalic 
acid, C e H 4 (COOH) 2 . These materials condense together when heated, 
eliminating molecules of water between them. This is exactly similar 
in principle to the production of ethyl acetate (above) from acetic 
acid and ethanol, though more complex. The first stage is: 
2C 2 H 4 (OH) 2 + C,H 4 (COOH) 2 -»- C 8 H 4 (C s H 4 2 ) 2 (OH) t + 2H,0 

The organic product is still an alcohol (notice the two OH groups) 
and it can condense with more acid, a process which continues until 
the molecular weight of the product, terylene, exceeds 50 000. Tery- 
lene can be extruded as fibres and woven into fabrics, sometimes in 



combination with other textiles. It is not a true polymer (p. 329) 
because water is eliminated from the materials as it is synthesised, 
but, because of its complexity, it is known as a polyester. 


Acetic acid is the second member of the homologous series of fatty 
acids, so called because some members of the series, e.g., stearic acid, 
C 17 H 35 .COOH, can be prepared from fats which are their esters. The 
general molecular formula of the acids is CnH^^.COOH, or 
CHOC'S. The first formula is usually preferred because it expresses 
the presence of the carboxyl group, COOH, which is the characteristic 
group of the series. The first two members of the series are formic acid 
H.COOH or CH 2 2 , and acetic acid, CH 3 .COOH or C a H 4 2 . Notice 
that, as in all homologous series, successive members have molecular 
formula? which differ by CH 2 . 

Laboratory preparation of acetic acid 

Acetic acid is the final oxidation product of ethanol. The oxidation 
occurs in two stages-the first to aldehyde, the second to acetic acid 
Both stages can be carried out in the laboratory by heating with 
chromic acid (i.e., potassium dichromate in dilute sulphuric acid) 
CH 3 .CH 2 OH -► CH 3 .CHO + 2H+ + 2e~ 
CH3.CHO + H 2 -> CH3.COOH + 2H+ + 2e~ 
In practice, this method gives a dilute solution of the acid from which 
a pure product is not easily obtained. An easier preparation is the 
distillation of anhydrous sodium acetate with concentrated sulphuric 
acid. Hydrochloric acid is not suitable because it is volatile and would 
distil with the acetic acid. 

CH s .COONa f H 2 S0 4 -»- NaHS0 4 + CH 3 .COOH 
The apparatus must be all-glass because hot acetic acid vapour attacks 
cork or rubber stoppers. A retort is suitable (see nitric acid p 435) 
or a modern glass-jointed distillation apparatus. Acetic acid distils 
as a colourless liquid. 

Large-scale production of acetic acid 

Pure acetic acid is now made on the large scale from acetylene 
C 2 H 2 , in two stages. In the first stage, acetylene is passed into dilute 
sulphunc acid at about 96°C, with mercury(/I) sulphate present as 

C a H 2 + H 2 0-*CH 3 .CHO 

The aldehyde vapour produced is then passed over a heated catalyst 
manganeseiJI) acetate, with air as oxidising agent. 
CH a .CHO + ±0 2 -»- CH3.COOH 




Imperial Chemical Industries U4. 

P.- ate 8. (a) Top section of .he Oslo evaporator for making granular salt, in .he I.C I Works 

at Slokc. 

PLA *JL ( . h) An opcn : he:,r r [h furnace . being charged with raw material for steel-makinc The funwt 
shown has a capac.ty of .53 tons, taking about nine hours for melting and relining 'he charg.T 

L^ a^^^^^^^^^— ^^— ^— ,-,. Slemirl\ u-ul I /nl'ilt Ltd. 

Acetic acid is also manufactured in the form of vinegar, which 
contains about 4% of the acid. A dilute aqueous solution of ethanol, 
e.g., a wine of poor quality, is run slowly over coke in a well-aerated 
vat. The coke is smeared with 'mother of vinegar', a white solid 
material containing the bacterium, mycoderma aceti. This acts as a 
catalyst to oxidation of ethanol by oxygen of the air. The liquid may 
have to pass through the vat several times to secure full oxidation. 

CHa-CHgOH + 2 ->CH 3 .COOH + H 2 

Properties and reactions of acetic acid 

Physical properties. In English conditions, acetic acid is usually a 
colourless liquid with a strong, very irritating smell. It mixes with 
water in all proportions and a dilute solution has the usual sour taste 
of an acid. The melting-point of the acid (17°C) is near room tem- 
perature. In winter, the pure acid may freeze to crystals resembling 
ice; hence the name glacial acetic acid which is usually used for the 
pure material. 

Chemical reactions. (1) Acidic behaviour. Acetic acid is a weak acid, 
that is, it is slightly ionised in dilute solution. For example, at 0.1M, 
or 6 g/dm 3 , the acid has a degree of dissociation of 1.4%; at M, it is 

CH3.COOH ^ CH3.COO- + H + 

The presence of H + gives the solution the usual acidic behaviour, as 

It turns blue litmus to red. 

It forms a salt with an alkali or base. 

CH3.COO- + H + + Na+ + OH--> CH 3 .COO-Na+ + H s O 
It liberates carbon dioxide from a carbonate or bicarbonate. 

CO s a - + 2H+ -> H 2 + C0 2 

HCO3- + H+ ->- H 2 + C0 2 

It liberates hydrogen on contact with a strongly electropositive 
metal, e.g., Mg. 

Mg + 2H+ ->• Mg 2+ + H 2 
All compounds containing the carboxyl group, COOH, show these 
weakly acidic reactions. This includes all the fatty acids and also 
compounds which do not belong to this series, e.g., oxalic acid and 
tartaric acid. These are both dicarboxylic acids, containing two car- 
boxyl groups per molecule. 

Oxalic acid Tartaric acid 

H 8 C 2 4 or (COOH) 2 C 2 H 4 O s (COOH) 2 

Both acids can neutralise alkali in two stages, giving an acid salt and 





a normal salt. For oxalic acid with sodium hydroxide solution, the 
stages are: 

H 2 C 2 4 + NaOH -► NaHC,0 4 + H,0 

H 2 C 2 4 + 2NaOH -> Na 2 C 2 4 + 2H 2 

The other acidic properties given for acetic acid are also shown by 

these acids, but the liberation of hydrogen with metal may be negligible 

for very weak organic acids. 

(2) Ester formation. A mixture of glacial acetic acid and moderate 
excess of a simple alcohol, e.g., CH 3 OH or C„H s OH, boiled under 
reflux with a little concentrated sulphuric acid or dry hydrogen chlor- 
ide present as catalyst, produces an ester. For example, with ethanol : 

CH 3 .COOH + C 2 H 6 OH ^ CH 8 .COOC 2 H 6 + H 2 
For more details, see p. 336. 

(3) With chlorides of phosphorus. Glacial acetic acid reacts rapidly 
in the cold with phosphorus pentachloride. 'Steamy' fumes of 
hydrogen chloride are given off and the organic product is acetyl 
chloride, a colourless liquid which fumes in air. 

CH s .COOH + PCI 6 -»- CH 3 .COCl + HC1 + POC1, 

This reaction shows the presence of the hydroxyl group, OH, in the 
acid; it is part of the carboxyl group, COOH. 

Phosphorus trichloride reacts similarly but more slowly: slight 
heating is required. 

3CH s .COOH + 2PCI 3 ->- 3CH 3 .COCl + P 2 O s + 3HC1 

Notice that the chlorides of phosphorus attack the COOH group of 
the acid but leave the CH 3 group unchanged. 

(4) With chlorine. If glacial acetic acid is boiled under reflux while 
chlorine is passed in, monochloroacetic acid is produced. Light 
catalyses this reaction (photocatalysis). 

Cl 9 + CH3.COOH -► CH,Cl.COOH + HC1 
Further passage of chlorine will give, in succession, di- and trichloro- 
acetic acid, CHCl 2 .COOH and CCl 3 .COOH respectively. Notice that 
chlorine attacks only the CH 3 group, leaving the COOH group intact. 
If monochloroacetic acid is allowed to react with aqueous am- 
monia, it is converted to aminoacetic acid, also called glycine. NH t is 
called the amino group. 

2NH 3 + CH 2 Cl.COOH ->■ CH a NH 2 .COOH + NH 4 C1 
Excess of ammonia would produce the ammonium salt of glycine 
CH 2 NH 2 .COONH 4 . 

Glycine and other, more complex, aminoacids are very important 
biological materials. For example, a protein, boiled with dilute acid, 
e.g., H 2 S0 4 , is hydrolysed into aminoacids, such as glycine and 

alanine, CH 3 .CH(NH 2 ).COOH. This occurs because the protein 
contains peptide linkages, -CO.NH— , and they break down in 
hydrolysis by using the elements of a molecule of water. 
-CO.NH— + H s O->- — COOH + NH 2 — 
The COOH and NH 2 groups are then constituents of different amino- 
acids. Quite complex mixtures of aminoacids and other organic 
mixtures of compounds can now be investigated very efficiently by 
the technique of chromatography, about which something will be 
written in the next section of this chapter. 

Complex compounds closely resembling natural proteins can be 
built up in the laboratory by condensing aminoacids together, i.e., 
causing them to combine by eliminating water molecules between the 
NH S group of one aminoacid and the COOH group of another. This 
forms a peptide linkage in the course of creating a more complex 

R.COOH + NH 2 .R, -> R.CO.NH.R, + H 2 
R and K t represent the rest of the aminoacid molecules. The con- 
densation can be continued to give a quite large, protein-like mole- 
cule, but with nothing like the complexity of a natural protein so far. 


Chromatography is a means of separating the constituents (usually 
organic) of a mixture by taking advantage of their different rates of 
movement (in a solvent) over an absorbent medium. For each con- 
stituent, the rate of movement depends on the relative affinities of 
the constituent for the solvent and the absorbent medium. In paper 
chromatography, which is one of its most useful forms, the absorbent 
medium is filter-paper and a variety of solvents can be used, e.g., 
ethyl acetate, methanol, aqueous phenol. The following experiment 
illustrates the process of chromatography in a simple way. A pin- 
hole is made at the centre of a Whatman No. 1 filter-paper of 12.5 cm 
diameter. String is threaded through the hole and cut to 0.5 cm 
length on one side and 5 cm on the other. A drop of BDH universal 
indicator is then placed in the centre of the filter-paper and allowed 
to dry out and turn brown in air. A Petri dish (or shallow crystallising 
dish) of 10 cm diameter is about one-third filled with the solvent, 
ethyl acetate, and the filter-paper is placed over it so that the longer 
end of string dips into the solvent. Another, similar dish is placed on 
top (Fig. 1 1 1), to make a closed system. Ethyl acetate will ascend the 
string and spread over the filter-paper so developing a chromatogram 
of the constituents of the indicator. That is, after a suitable time (up 
to an hour), bands of the different dyes present will be seen— a purple 
colour ahead, a yellow colour about halfway and a red-brown dye 



remaining close to the centre. If methanol is the solvent, an outer 
yellow band occurs, with an inner red band and no residue at the 

In standard chromatography, rectangles of filter-paper about 22 
cm by 5 cm are usually used. A spot of solution containing the 
mixture under investigation is placed near the middle of one of the 
short ends of the filter-paper and about 3 cm away from the end. 
Then the spot is allowed to dry. The paper is then suspended verti- 
cally with the spotted end immersed in the solvent but the spot well 
clear of the liquid level. The whole is enclosed by a covering vessel 
to prevent loss of solvent. The chromatogram will then develop as the 
solvent rises up the filter-paper. Spots of the various constituents of 

Dried spot of indicator (BDH universal) 

Petri dish 
{or similar) 

ethyl acetate 
(or other solvent) 
Fio. 111. 
A simple example of chromatography. 

the original mixture will collect, at different distances along a verti- 
cal line, above the original spot. When the ascending front of solvent 
is approaching the top of the paper (usually eight hours or more 
after the start), the paper is removed and dried. It can then be ex- 
amined in various ways to locate the spots of material left and to 
identify their contents. 

For example, the paper may be examined in ultra-violet radiation, 
which may cause fluorescence and so indicate the position of some 
compounds; or the paper may be sprayed with a reagent giving colour 
reactions with compounds present. If desired, areas of paper on 
which spots of materials have been detected may be cut out and 
'eluted' with a solvent, so providing a solution in which a compound 
may be identified by the usual methods of micro-analysis. Two of the 
advantages of chromatography are that it can be performed with very 
small weights of material and they are not destroyed in the process. 



To secure better separations, the following device may be used. A 
square piece of filter-paper is employed and a chromatogram is 
developed from a spot of material near one corner of the paper. This 
gives spots of separated materials along a line, as above. The paper 
is then dried and a second chromatogram is developed by using the 
same piece of paper at right angles to its former position and with a 
different solvent. Spots of material will then be found, after drying, 
much more widely spread over most of the surface of the paper. 


Amines are members of the homologous series which has a general 
molecular formula of C„H M , ,NH 2 for its constituents. The first two 
members (both gases in ordinary conditions) are methylamine, 
CH 3 NH 2 , and ethylamine, C a H 6 NH 2 . The amino group, NH 2 , is the 
characteristic group of all members of the series. Like all such amines, 
these two compounds are derived from ammonia by replacing one 
hydrogen atom of the NH 3 molecule by an organic group such as 
CH S or CjH 8 . 

H C 2 H 5 

H— N-H 


H—N-H or C 2 H 6 NH 2 

Properties of simple amines. Methylamine and ethylamine are 
colourless gases at room temperature and pressure; they have a 
strong smell of fish and ammonia and, like ammonia, are very 
soluble in cold water. They burn in air and this distinguishes them 
from ammonia, which does not. 

Basic nature of amines 

Like ammonia, the amines given an alkaline reaction in aqueous 
solution by producing a considerable concentration of hydroxide 
ion, OH-. 

C 2 H 6 NH 2 + H 2 ^ C 2 H B NH S + + OH" 
Compare NH 3 + H s O «* NH, + + OH" 

Both methylamine and ethylamine turn damp red litmus paper blue 
and are, in fact, stronger bases than ammonia at corresponding 


Just as ammonia forms salts with mineral acids, the amines pro- 
duce corresponding salts, which can be obtained as similar white 

C 2 H 6 NH, + HC1 -> C,H 5 NH 3 + C1- (ethylamine chloride) 






NH S + HC1 -»• NH/C1- (ammonium chloride) 
These amine salts liberate the free amine when warmed with aqueous 
caustic alkali, just as an ammonium salt liberates ammonia in the 
same conditions. 

QH 5 NH 3 +C1- + Na + OH- -»- Na + Cl" + H a O + C,H S NH 3 

NH«+C1- + Na+OH- -»■ Na+Cl" + H,0 + NH 8 

Like aqueous ammonia, a solution of a simple amine supplies a large 
enough concentration of hydroxide ion to precipitate the insoluble 
hydroxides of several metals from solutions of their salts, e.g., iron(lll), 
zinc and aluminium hydroxides. 

Fe 8 + + 30H- -► Fe(OH) 8 (Al similar) 
Zn 2 + + 20H- -»- Zn(OH), 

Preparation of an amine 

Ethylaminc can be prepared by the reaction between ethyl iodide 
(a liquid) and concentrated aqueous ammonia. Heating at 100°C in 
a sealed tube is required. 

C.H.I + 2NH 3 ->- C 3 H 5 NH a + NH 4 I 
By-products are also produced and the isolation of the amine is 
difficult. Methylamine can be prepared by a corresponding method 
from methyl iodide, CH 3 1. 


Nylon is usually called a superpolymer though it is not a true poly- 
mer (p. 329) because water is split out between pairs of molecules as it 
forms. These molecules are of the types HO — X — OH and H — Y— H 
and the essential reaction is: 

. . .- |H + HOI— X ^OH + Hr -Y- JH + HOl -X— 

OH + H-Y— H + HO— X— OH + H — ... 

— X— Y— X— Y— X— Y— X— Y— + mH,0 

This produces a greatly extended linear molecule which is nylon. 
The process is done in two stages, the second with heat and pressure 
in an autoclave and, after cooling, the nylon appears as chips. By 
melting the chips and forcing the liquid through tiny holes in a metal 
disc (a spinneret), filaments are formed, stretched between rollers 
and then gathered into nylon yarn. This yarn can be woven into a 

and H— NH(CH 2 ) 8 NH— H 


variety of attractive garments (underwear, stockings, shirts) and into 
tow ropes, tyre-cord, nets and racquet strings, which require strength 
and durability. Nylon powder can be moulded, while hot, into 
various plastic articles, e.g., gears and bearings in small machines 
such as electric mixers and razors, where toughness is required. 

(The actual compounds used in making nylon, both petroleum 
derivatives, are: 

HO— CO(CH 3 ),CO-OH 
adipic acid 

From this, X is -CO(CH 3 ) 4 CO— and Y is — NH(CH,),NH— and 
the linear nylon molecule is: 

. . . — CO(CH g )«CO— NH(CH 2 ) 8 NH— CO(CH 1 ) 1 — 

CO— NH(CH 2 )„NH- . . . 

Notice the peptide linkages, — CO— NH— , in this chain. It is found 
that, in 'condensation polymers' like nylon, the total number of 
CH 3 groups in the two constituents must be at least eight, and pre- 
ferably nine. Otherwise, cyclic compounds of small molecular 
weight will form, not very large linear molecules as in nylon.) Two 
other important linear polymers are the following. 


Styrene is phenyl ethylene, C„H 6 .CH=CH 3 . It is a colourless liquid 
prepared (in two stages) from benzene, C B H„, and ethylene. It poly- 
merises slowly at ordinary temperature and rapidly if heated to 
100-150°C, preferably in nitrogen, to prevent oxidation. It produces 
polymers of average molecular weight about 120 000 at 100°C and 
about 23 000 at 150°C. The polymers have a linear structure resembl- 
ing that of polythene (p. 329), but with a phenyl (C 8 H 5 ) group attached 
to each alternate carbon atom. 

— CH t — CH— CH,— CH— CH,— CH— CH 3 -CH— CH a -CH— 
C.H. C.H 5 C a H B C 8 H 5 C 8 H 6 

The polymer is a white solid and can be depolymerised (back to 
styrene) at about 350°C. 

Polystyrene is a very good electrical insulator (especially for high- 
frequency work) and, being water-resistant, can be moulded into 
battery cases and refrigerator parts. Expanded polystyrene (in the 
form of blocks and mouldings) can be used in thermal insulation 
and also in packaging and display. Great quantities of styrene 
(25%) are co-polymerised with butadiene (75%) in making a 
synthetic rubber. 





Perspex (methyl methacrylate polymer) 

An acidic derivative of ethylene is acrylic acid, CH s =CH.COOH. 
Methylacrylic acid is CIV=C(CH 3 ).COOH and its methyl ester is 
CH s =C(CH 3 ).COOCH 3 . This ester, methyl methacrylate, having a 
double bond like ethylene, polymerises in a corresponding way. The 
polymerisation is carried out at 80-90°C in an emulsion of the ester 
in dilute aqueous soap solution with hydrogen peroxide present. 
After cooling, the mixture is poured into very dilute hydrochloric 
acid when the polymer coagulates and is separated, washed and 
dried. The polymer is linear, like polythene, but has CH 3 and 
COOCH 3 groups attached to alternate carbon atoms. 



I 2 — C— CH, 





The polymer is colourless and transparent. When softened by heat, 
it can be moulded into objects with optical uses, e.g., lenses and 
reflectors, and, in sheet form, as 'Perspex' can replace the much 
denser glass in aeroplane domes and windows. It is also moulded 
into artificial dentures and small household articles. (Notice how, in 
both polystyrene and the methyl methacrylate polymer, the double 
bonds of the simple compound are lost and the polymer is a saturated 


1. What is meant by the term hydrocarbon! Give the molecular formula 
of one gaseous alkane and one gaseous alkene hydrocarbon. Contrast 
the reactions of these two with chlorine. Give one good test by which the 
two gases may be distinguished and explain it. 

2. State briefly without diagrams how ethylene and acetylene are pre- 
pared and collected in the laboratory, mentioning suitable conditions of 
reaction. Contrast the behaviour of these gases with pure chlorine. Give 
the chemistry of the large-scale preparation of one important product from 
each of these hydrocarbons and mention how the product is used in each 

3. QHjOH is the molecular formula of the second member of the 
homologous series of alcohols. Write the molecular formula of the fourth 
member. Mention two differences of physical properties which you would 
expect the two compounds to show. Give two examples of similarity of 
chemical behaviour you would expect from these compounds. 

4. What is a peptide linkage? How does a peptide linkage of a protein 
behave when the protein is hydrolysed? What are the resulting products? 
Explain what value paper chromatography has in investigating reactions 
of this kind and state, in outline, how the experimental work is carried out 

5. Show, by equations and brief comment, the chemical relation between 
ethanol and acetic acid. How, and in what conditions, does acetic acid 
react with (a) chlorine; (b) phosphorus pentachloridc ? Describe briefly the 
preparation (without purification) of one ester of acetic acid. How does 
this ester behave if boiled with aqueous sodium hydroxide? 

6. Explain why acetic acid shows 'acidic' reactions and state two such 
reactions (other than behaviour with litmus). In what way, and why, does 
the acidity of this acid differ from that of hydrochloric acid ? State (a) in 
outline a good method for the preparation in the laboratory of reasonably 
pure acetic acid; (b) the chemistry of one large-scale method of making 
acetic acid. 

7. Explain the basic nature of ethylamine, C,H 6 NH„ and compare the 
behaviour of this compound with that of ammonia towards hydrochloric 
acid. How, by one chemical test, would you distinguish ethylamine from 
ammonia (both gaseous)? If ethylamine is the second member, write the 
molecular formula of the fifth member of the homologous series. Mention 
two differences of physical properties you would expect between the two 

8. How, by one chemical test in each case, would you distinguish between : 
(a) ethane and ethylene; (b) ethanol and acetic acid; (c) ethylene and acety- 
lene? Outline one conversion of ethylene into ethanol. 

9. Explain what is meant by polymerisation. Outline the manufacture 
of polythene. Write a short account of its uses and why it is suited to them. 
Explain why terylene is called a polyester. Why is it not a true polymer? 
What is the chief use of terylene in ordinary life? 

10. There are two isomers of molecular formula, C,H»0, i.e., an alcohol 
and a compound of another type. Write their full structural formula;. 
Give one chemical test to distinguish between them. Write full structural 
formula; for the two isomers of butane, C 4 H 10 . Pentane, QH,,, has three 
isomers. Write full structural formula: for two of them. 



Chapter 25 

Silicon and its Compounds 

(For silicon in periodicity, see p. 87) 

SILICON is the second most abundant element in the earth's crust, 
the most abundant of all being oxygen. Silicon is found in the 
following forms: 

1. As metallic silicates. Igneous rocks, such as granite and basalt, 
consist largely of mixtures of silicates, those of magnesium, alu- 
minium, potassium and iron being most common. China clay, kaolin, 
is a hydrated silicate of aluminium; ordinary clay is a mixture of 
particles of quartz, mica and other substances bound together by a 
sticky material which is a hydrated silicate of aluminium of approx- 
imate formula Al 2 Si 2 7 .2H a O. 

2. As silica {silicon dioxide), SiO t . The purest form of silica is rock 
crystal or transparent quartz. Sand usually consists of silica, with 
various impurities, and opal, hornstone and jasper are forms of this 
oxide. It also exists in a less pure form as flint. 

The porous material 'kieselguhr', is a form of silica made up of the 
fossil shells of small plants called diatoms. It is used to absorb the 
explosive, 'nitroglycerine', forming the product called 'dynamite', 
which is safer to transport and handle than the explosive itself. 


The element may be prepared in the amorphous state by heating 
magnesium powder with silica, SiO a . 

SiO a + 2Mg->2MgO + Si 
The reaction is very violent. 

Silicon from silica 

Mix one 'saltspoonful' of magnesium powder with three times its 
bulk of amorphous silica and place the mixture in a test-tube. 
Clamp the tube with the mouth pointing upwards and in a safe direc- 
tion, place a Bunsen burner underneath the tube to heat it stronelv 


and retire to a distance of two metres or more. After a few moments 
the reaction occurs with a flash of light. Allow to cool, break the 
test-tube in a mortar and pick out the pieces of brown amorphous 

2Mg + Si0 2 -»• 2MgO + Si 

Reactions of silicon 

Place a small piece of silicon, obtained in the above experiment, on 
a tripod stand and direct the flame of a Bunsen burner on to it from 
above. It will react, sometimes by burning, to form a white solid, 


Si + 2 -*-SiO» 

Put the rest of the silicon into a test-tube, add sodium hydroxide 
solution and warm. There is an evolution of hydrogen and finally the 
silicon dissolves. 

2NaOH + Si + H 2 -> Na 2 SiO s + 2H 2 
Amorphous silicon combines with chlorine at low red heat and 
also decomposes steam. 

Si(c) + 2Cl 2 (g)->SiCl,(l) 
Si(c) + 2H 2 0(g)-»- Si0 2 (c) + 2H 2 (g) 

Silicon is used in the manufacture of certain types of steel and 


Silica from sodium silicate 

Add concentrated hydrochloric acid to a solution of sodium silicate 
and warm if necessary. White hydrated silica (silicic acid) will come 
down as a gelatinous precipitate. 

Na 2 SiO s + 2HC1 -»- 2NaCl + Si0 2 .H a O 
Show that the precipitate is readily soluble in dilute caustic soda 

2NaOH + SiO a -»- Na 2 SiO s + H 2 

Uses and properties 

In its naturally occurring form, sand, silica finds great use as a 
constituent of mortar and cement, for the manufacture of glass (see 
p. 352) and for filtration of water in bulk. 

Fused silica is used for the manufacture of certain types of labora- 
tory apparatus. Its coefficient of expansion is very low (about one- 
fiftieth of that of glass) and consequently the strains set up in it 
by irregular contraction under sudden reduction of temperature are 
small and insufficient to break it. A red-hot silica tube or crucible 



can be plunged into cold water without damage. Fused quartz 
threads are used in the construction of physical apparatus. They can 
be made so thin as to be invisible to the naked eye and yet strong 
enough to support a weight of 2 g. Their tenuity and elasticity 
make them invaluable for light, delicate suspensions. The difficulty 
in working fused silica is that the very high temperature of at least 
1500°C is required. 

Silicon dioxide (silica) is almost insoluble in water. It acts as the 
acidic oxide of a non-metal and forms silicates with caustic alkalis 
and metallic oxides (all heated). 

SiO, -f 20H- ->■ SiO s *- + H t O or SiO, + 0«- -»- Si0 3 2 - 

Silicon dioxide forms three-dimensional crystalline systems of great 
complexity. In its simplest form (cristobalite), the silicon atoms are 
arranged like the carbon atoms of diamond (p. 288) but with oxygen 
atoms midway between them. This macro formation accounts for 
the high melting point of silica in contrast with the gaseous state of 
carbon dioxide at s.t.p. This gas contains only separate molecules of 
carbon dioxide (p. 73). 

Sodium silicate from silica 

Wash some of the white gelatinous precipitate of hydrated silica 
obtained above and ignite it in a crucible. The product is amorphous 
silica and is no longer easily soluble in caustic soda solution even if 
hot and concentrated. That silica is an acidic oxide can, however, be 
shown as follows: Put a small piece of solid sodium hydroxide about 
half an inch long into a crucible, add a 'saltspoonfuP of amorphous 
silica and heat the mass strongly for about ten minutes. Take great 
care when doing this, that none of the molten mass is spilled or comes 
into contact with the flesh. Allow the melt to cool, fill the crucible 
two-thirds full of water and warm. The solution obtained is sodium 
silicate and from it gelatinous silica can be obtained by the action of 
hydrochloric acid. 

Silica is non-volatile and can, at high temperatures, displace 
volatile acidic oxides from combination. They pass off as vapour, 
leaving a silicate, e.g., 

Na,S0 4 + Si0 2 -»- Na 2 Si0 3 I SO a 

Silicates are salts of the silicic acids. The chemistry of these acids 
is complex, and little is known for certain about them. 

The gelatinous substance precipitated by acidifying a hot solution 
of 'water-glass' has the empirical formula H 2 Si0 3 , and is called 
meta-silicic acid, and an acid of approximate formula H 4 SiO«, 



ortho-silicic acid, is obtained by the action of water on silicon 

Na 2 Si0 3 + 2HC1 -*■ H 2 SiO„ + 2NaCl 
SiCl 4 + 4H s O -> H 4 Si0 4 + 4HC1 

The salts of these acids are called silicates. The meta-silicates are the 
most important. 

Sodium silicate (water-glass), Na a SK) 3 

Sodium silicate is prepared by heating two parts by weight of silica 
with one part of sodium carbonate. 

Na 2 C0 3 + SiO,-* NaaSi0 3 + CO, 

The product is a glassy solid. It is heated with water (under pres- 
sure) to dissolve it, and is sold in tins in the form of a concentrated 
solution, similar in consistency to 'golden syrup', but colourless. 
This is called 'water-glass'. 

Water-glass is chiefly used for preserving eggs. The shell of an egg 
consists largely of calcium carbonate. This, when the egg is im- 
mersed in a suitable solution of 'water-glass', reacts with the sodium 
silicate to produce a precipitate of calcium silicate which seals the 
pores of the shell, excluding bacteria and so preserving the egg from 

CaCOg + Na,Si0 8 -* CaSiO„ + Na 8 C0 8 


An interesting chemical phenomenon, called a 'silica-garden', can 
be shown in the following way. Dilute some water-glass until it has a 
density of 1.1 g/cm 3 (test with a hydrometer), then filter it and put 
the liquid into a tall, rather narrow vessel. Drop into it crystals of 
manganese(II) chloride, cobalt nitrate, iron(II) sulphate and 
copper(II) sulphate. From these crystals there will shoot fantastic 
coloured growths. Those from the cobalt salt usually appear within 
a few seconds and grow rapidly. They are dark blue. Growths from the 
other crystals appear more slowly. Those from the manganese salt 
arc pale pink, from copper sulphate light blue, and from the iron(II) 
salt green. The growths are tubes of the silicates of the metals. 

Silica gel 

If hydrochloric acid is added to a concentrated solution of sodium 
silicate at about 100°C, a product known as silica gel is precipitated. 
It contains 5-7% of water and is regarded as a partially hydrated 

Na 8 Si0 3 + 2HC1-* 2NaCl + SiQ 2 .H 2 



Silica gel absorbs water-vapour very readily and is used for drying 
gases on the industrial scale, e.g., in drying air for the blast of a blast- 
furnace. When spent, the gel can be re-activated by heating to a suit- 
able temperature. 

Silica gel can also absorb volatile solvents, e.g., carbon disulphide 
or acetone, which would otherwise be lost as effluent. The solvents 
can be recovered by suitable heating of the gel. It has also been used 
successfully for freeing petroleum oils from sulphur compounds. 


The substance to which the term 'glass' is usually applied consists 
of a mixture of two or more silicates. Common soft glass, of which 
bottles and laboratory apparatus are made, is prepared by heating 
together silica in the form of sand, sodium carbonate or sodium 
sulphate, and chalk or limestone (calcium carbonate). (Some broken 
glass and a little coke are usually added.) The glass so prepared con- 
sists of a mixture of sodium silicate and calcium silicate. 

NajCOg + SiQ 2 -*- Na 2 Si0 3 + C0 2 



Na 2 SO« S0 3 

CaC0 3 -f SiO a -»- CaSiOa + C0 2 

If potassium carbonate is used instead of sodium carbonate, a 
'hard' glass, that is, a glass needing a higher temperature to melt it, 
will be produced, and will consist of a mixture of calcium silicate and 
potassium silicate. 

Glass containing lead silicate has a brilliant appearance, and is 
made by adding to the ordinary glass mixture some red lead, Pb s 4 , 
or litharge, PbO. Such a glass ('flint glass') has a high refractive index 
and is used for prisms and lenses, but it is soft and should be wiped 
only with silk, to avoid scratching the surface. 

Coloured glass is obtained by addition of metallic oxides, for 
example, cobalt produces blue glass, chromium produces green glass. 
Opalescent glass may be produced by addition of calcium phosphate. 

Physical nature of glass 

Glass is not really a solid. This is shown by the fact that, when 
heated, it does not suddenly pass from the solid to the liquid state at 
a definite temperature, but softens slowly as the temperature rises 
and gradually becomes liquid. A true solid would melt suddenly over 
a small temperature range. Glass is a super-cooled liquid. By this we 
mean that its molecules have not taken up a definite formation to 
produce crystals, but are arranged at random as they were in liquid 


The softening of glass before it melts is a very valuable property. 
It enables a skilled worker or a machine to blow it into various 
shapes— bottles, flasks, beakers— by using suitable moulds. While 
soft, glass can also be moulded and joined to produce elaborate 
scientific apparatus. Large sheets of glass are produced by rolling 
out a mass of hot glass on a long, flat table or by floating glass on 
molten tin while heating the top surface. Glass-tubing is made by 
first blowing a mass of glass into a thick-walled cylinder, and then 
drawing the cylinder out between two workers walking in opposite 
directions, one manipulating the blow-iron. The 'unbreakable glass' 
used for motor-car windscreens is made by cementing a sheet of glass 
on to each side of a sheet of celluloid. The cement prevents splinters 
from flying off. 

Articles made of glass usually have to be annealed, that is, heated 
to a suitable temperature and then allowed to cool very slowly and 
uniformly so that no stresses are set up which might break the glass. 
The annealing of considerable masses of glass may require several 
weeks, or even months, as in the case of the glass for the mirror of 
the 500 cm telescope erected at the Mount Wilson observatory in the 
U.S.A. Working with glass is a very ancient art. It was practised by 
the Egyptians at least 2300 years ago, while some of the finest speci- 
mens of glassware are the work of Venetian craftsmen of the Middle 

Silicon tetrachloride 

Silicon combines with chlorine when heated to form silicon tetra- 
chloride, but a more convenient method of producing this compound 
(because it does not require the element, silicon) is to heat an intimate 
mixture of silica and carbon in a current of dry chlorine. 

SiO, + 2C + 2C1 2 -► SiCl< + 2CO 
Silicon tetrachloride distils off as vapour and can be condensed to a 
colourless liquid of boiling-point 58°C at 760 mm. It may contain a 
small proportion of other chlorides, e.g., Si 2 Cl e . 


(1) Silicon tetrachloride reacts rapidly with cold water, forming 
hydrogen chloride and silicic acid. 

SiCl, + 3H 2 -*■ 4HC1 + H 2 SiO a 
This reaction (with atmospheric moisture) causes the tetrachloride to 
fume in moist air. Carbon tetrachloride, CC1 4 , however, has no 
reaction with water, which is unusual for the chloride of a non-metal. 

(2) If heated with an alkali-metal, silicon tetrachloride is decom- 
posed to form silicon. 

SiCl« + 4K -»• Si + 4KC1 




1. What is the chemical nature of glass? How is it made? Why is glass 
stated to be a super-cooled liquid? 

2. Silica is an acid-forming oxide. Justify this statement by reference to 
the chemical properties of silica. 

3. In what forms docs silica occur in nature? How may the element 
silicon be obtained from silica? 

4. What is silica? Starting with the naturally occurring substance 
describe how you would obtain (a) water-glass; (Z>) a solution of silicic acid ; 
and (c) pure silica. (L.) 

5. Describe in outline two methods by which silicon may be prepared 
in the laboratory from silica, SiO,. How, and in what conditions, does 
silicon react with (a) sodium hydroxide, (6) chlorine, (c) oxygen? Compare 
the behaviour of carbon with that of silicon towards these three reagents. 

6. Describe a laboratory preparation of a concentrated solution of 
sodium silicate from silica, SiO t . How may the product silica gelbe obtained 
from this solution? Briefly discuss the large-scale uses of this product. 

7. Describe in brief outline a preparation of silicon tetrachloride from 
silica, SiOj. Compare and contrast the elements carbon and silicon with 
reference to the chemical nature of their dioxides and the behaviour of 
their tetrachlorides with water. Mention the experimental evidence for your 

Chapter 26 


THE four non-metals, fluorine, chlorine, bromine and iodine, 
make up a family of related elements, the chemical properties of 
which form an interesting study. The first of the series is the very 
active element fluorine, which is so active that it evaded efforts to 
isolate it for many years, because it reacted with almost every element 
or substance with which it came into contact— even the glass of the 
apparatus in which the reaction took place! Of fluorine we need say 
but little since it is so difficult to isolate that its properties could have 
no practical application in the school laboratory, and of the others 
we shall see that bromine is almost an 'arithmetical mean' between 
chlorine and iodine. Chlorine was first isolated by Schcele in 1774. 

Scheele was a Swedish apothecary' who carried out, during the 
short time which he lived (he died when only 44 years of age), a vast 
number of illuminating experiments, performed in an old shed 
attached to his house. It was whilst he was investigating the properties 
of pyrolusite (impure manganese dioxide) that he heated it with con- 
centrated hydrochloric acid in a retort and collected the chlorine 
gas which came off into a pig's bladder. 

It is usually made by the oxidation of concentrated hydrochloric 
acid. The oxidation can be brought about by many oxidising agents, 
for example, lead dioxide, manganese dioxide, red lead or potassium 

o + 


, i 

oxygen from 
oxidising agent 




H s O + Cl g 

water chlorine 

Preparation of chlorine from concentrated hydrochloric acid by oxi- 
dation with manganese dioxide (manganese(IV) oxide) 
Fit up the apparatus as shown in Fig. 112. Put some manganese 

dioxide into the flask, pour concentrated hydrochloric acid down the 






funnel and shake well before connecting up the flask with the rest of 
the apparatus. (Note. The use of a gas-ring and gauze keeps the 
flask low and makes the apparatus more stable.) 

Heat the mixture in the flask and effervescence is observed. A 
greenish-yellow gas is evolved which, together with a certain amount 
of hydrogen chloride (misty fumes), passes over into the first bottle 

Water to remove 
Hydrogen Chloride 

Fio. 112. 
Preparation of chlorine. 

*Conc Sulphuric 
acid to dry the. gas 


which contains water. This removes the hydrochloric acid gas (which 
is very soluble in water), and the concentrated sulphuric acid in the 
second bottle dries the gas which is collected by downward delivery, 
the gas being denser than air. 

MnO,(c) + 4HCl(aq) -► MnCl, + 2H,0 + Cl*(g) 
The above experiment should be carried out in a fume-chamber, as 
should any preparation of chlorine in which it is collected by dis- 
placement of air. 

Preparation of chlorine from concentrated hydrochloric acid by 
oxidation with potassium permanganate 

[If the chlorine is required pure and dry, insert wash-bottles containing 
(a) water, and (6) concentrated sulphuric acid in Fig. 113 and collect by 
displacement of air.] 

2KMnO t + 16HC1 -> 2KC1 + 2MnCl, + 8H s O + 5C1 8 
This is a very convenient laboratory method because it takes place 
in the cold, and, if the gas is collected over brine, the experiment need 
not be conducted in a fume-chamber. 

Solid potassium permanganate is placed in a flask and concen- 
trated hydrochloric acid is dropped on to it from a tap-funnel (Fig. 
113). As each drop of acid reaches the permanganate, there is evolved 

at once the corresponding quantity of chlorine. The apparatus is filled 
with the greenish-yellow fumes of the gas. Several gas-jars of the gas 
should be collected and the experiments described later performed to 
illustrate its properties. 

' Hydrochloric 


Fig. 113. 
Preparation of chlorine. 

Preparation of chlorine from common salt (sodium chloride) 

Chemistry of the action. Concentrated sulphuric acid acts upon 
the sodium chloride to form hydrogen chloride, which is then oxi- 
dised to chlorine by manganese dioxide. 

The apparatus is identical with Fig. 112, and the experiment is 
performed in a similar manner except that an intimate mixture of 
sodium chloride and manganese dioxide is placed in the flask and con- 
centrated sulphuric acid added. In this experiment the presence of 
hydrogen chloride as an impurity is more obvious. It is removed by 
passing the gases through water, and the chlorine is dried by means 
of concentrated sulphuric acid and collected as shown in Fig. 112. 

2NaCl + MnO t + 2H S S0 4 -* Na 8 SO« + MnS0 4 
common manganese concentrated sodium manganese 
salt (sodium dioxide sulphuric sulphate sulphate 

chloride) mangancsc(IV) acid 

oxide + 2H 2 -+- Cl t 

water chlorine 

Preparation of chlorine from bleaching powder 

In this case the chlorine is not prepared by the oxidation of con- 
centrated hydrochloric acid. The apparatus used is identical with 
Fig. 113 (above), bleaching powder is placed in the flask and a 
dilute acid, e.g., nitric acid, is dropped on to the powder. Efferves- 
cence occurs and the greenish-yellow gas can be collected by either 
of the methods mentioned previously. Heat is not required. 




CaOCl 2 + 2HN0 3 - 
bleaching dilute 

powder nitric acid 

• Ca(N0 3 ), + H,0 + CI 2 
calcium water chlorine 


CaOCl, + 2HC1 -+ CaCl 2 + H,0 + CI, 
hydrochloric acid 

CaOCl, + H^O,-*- CaSO« + H 2 + CI, 
sulphuric acid 

Industrial preparation 

By electrolysis of brine (see p. 253). 
Test for chlorine. Chlorine is a greenish-yellow gas which rapidly 
bleaches damp litmus paper. 

Isotopes of chlorine 

Chlorine has two principal isotopes. Both have the electron ar- 
rangement 2, 8, 7. One has an atomic nucleus of 17 protons and 18 
neutrons (CI = 35); the other has a nucleus of 17 protons and 20 
neutrons (CI = 37). Ordinary chlorine (CI = 35.5) is a mixture of 
these isotopes in the proportion of three of the lighter to one of the 
heavier atoms. 

Properties of chlorine 

Chlorine is a greenish-yellow gas with a choking, unpleasant, 
irritating smell. It is very poisonous if inhaled to even a small extent 
(1 part of chlorine in 50 000 of air may be injurious). It was used 
extensively during the war of 1914-18 and, being about 2± times as 
dense as air, it would roll along the ground when propelled by a very 
gentle wind without a great deal of it escaping upwards. 

It bleaches damp litmus and is a very reactive gas indeed. The fol- 
lowing experiments illustrate its properties and they are classified 
according to the various ways in which the gas can act. 

Chlorine as a bleaching agent 

Pour a little litmus solution into a gas-jar of the gas. The litmus 
immediately turns colourless. Chlorine will bleach the colour from 
most dyes and will remove writing ink (but not printer's ink, which 
consists mainly of carbon, which chlorine does not attack). 

Bleaching action. The chlorine reacts with the water, forming hypo- 
chlorous acid. 

Cl 8 + H 2 

-> HOC1 + HC1 

chlorine water 

hypochlorous hydrochloric 

acid acid 



This hypochlorous acid is a very reactive compound and readily 
gives up its oxygen to the dye, to form a colourless compound, 
dye + HOC1 -»- HC1 + (dye -I- O) 
coloured colourless 

Notice that hydrochloric acid is produced whenever chlorine 
bleaches, and hence an article must be thoroughly washed after 
bleaching or it will be attacked by the free acid. 

In industry, the article to be bleached is dipped into a tank con- 
taining bleaching powder in water, then into very dilute sulphuric 
acid. It is then washed by water to remove acid and may be treated 
with an anti-chlor to remove remaining chlorine, which, if left, might 
rot the material. A typical anti-chlor is sodium thiosulphate ('hypo'), 
which reacts with chlorine: 

S 20aS - + 4C1, + 5H,0-* 2S0 4 2 - + 8C1- + 10H+ 

Dry chlorine will not bleach. Pour about 20 cm 3 of concentrated 
sulphuric acid into a gas-jar of the gas as soon as it is collected. After 
a time put a piece of dry coloured cloth into the gas-jar. At the same 
time put a piece of damp coloured cloth into a gas-jar of chlorine; 
the latter is immediately bleached, whereas the former remains un- 
attacked. (If the chlorine and cloth are perfectly dry, the cloth remains 
unbleached indefinitely.) 

From the equation above, it is seen that water is necessary 
for the formation of hypochlorous acid, the compound 
which liberates the oxygen and whicli performs the bleach- 
ing. Hence, if no water is present no bleaching can occur. 

Chlorine as an element which combines readily with hydrogen to form 
hydrogen chloride 

(Bleaching can also be considered under this heading) 
Exposure of chlorine water to sunlight. Pass chlorine gas into water 
in a beaker for some time until the water becomes quite yellowish 
green in colour. Fill a long tube with this chlorine water, invert it in 
a beaker containing some of the water and expose to bright sunlight 
(Fig. 1 14). After some time, a gas collects in the tube and on applying 
a glowing splint, the latter is rekindled, showing the gas to be oxygen. 

2C1- + 2H,0 -*• 4HC1 + 2 
The chlorine has combined with the hydrogen of the water to form 
hydrogen chloride and oxygen is liberated. 

The above reaction probably occurs in two stages, as indicated by 
the equations: 

H,0-f Cl 2 -»- HOC1 + HC1 (instantaneous) 
water chlorine hypochlorous hydrochloric 





2HOC1 -»- 2HC1 + O t (slow) 
The burning of hydrogen in chlorine. Lower a jet of burning hydro- 
gen (great care before lighting, see p. 269) into a gas-jar full of 
chlorine. The hydrogen continues to burn with a white flame and 
clouds of steamy fumes of hydrogen chloride are seen, whilst the 
yellowish-green colour of the chlorine gradually disappears. 

H, + CI, -► 2HCI 

hydrogen chlorine hydrogen 

r . Oxygen 


. C/i/orme. 

Fio. 114. 
Exposure of chlorine water to sunlight. 

The readiness with which hydrogen and chlorine combine together is 
so great that, if a tube containing equal volumes of chlorine and 
hydrogen is exposed to sunlight, it explodes. 

Action of chlorine and warm turpentine. Warm a little turpentine in 
a dish, dip into it a filter-paper, and then drop this into a gas-jar of 
chlorine. There is a red flash accompanied by a violent action whilst 
a black cloud of solid particles of carbon is also formed. Hydrogen 
chloride can be shown to be present by blowing the fumes from an 
ammonia bottle across the top of the jar, when dense white fumes of 
ammonium chloride are observed. 

NH 3 (g) + HCl(g)-> NH 4 + CI-(c) 



Turpentine (a hydrocarbon) consists of hydrogen and carbon in 
chemical union. The chlorine combines with the hydrogen and leaves 
the black carbon behind. 

C 10 H l9 0) + 8Cl g (g) -+ 10C(c) + 16HCl(g) 
Effect of chlorine on a burning taper. Lower a burning taper into a 
gas-jar of chlorine. It burns with a small, red and sooty flame. Wax 
consists mainly of hydrocarbons and, as with the turpentine, the 
hydrogen forms hydrogen chloride and leaves the carbon. 

Action of chlorine on hydrogen sulphide. Invert a gas-jar of hydro- 
gen sulphide over a gas-jar of chlorine and remove the plates. You 
will observe a yellow precipitate of sulphur, and hydrogen chloride 
will be formed. 

H 3 S + CU -+ 2HCI + S 
hydrogen chlorine hydrogen sulphur 
sulphide chloride 

This can also be regarded as a case of oxidation-reduction, in which 
chlorine is reduced to its ions by electron gain and the sulphide ion 
(from H a S) is oxidised to sulphur by electron loss. Chlorine is the 
oxidising agent, hydrogen sulphide the reducing agent. 
[ H 2 S ^ 2H+ + S»- 
[S*- — > S + 2e~ (oxidation) 
Cl s + 2e~ -*■ 2C1~ (reduction) 


Chlorine as a chloride former 

Chlorine is a very reactive element and will combine with most 
other elements to form chlorides. In many cases chlorine will com- 
bine with elements spontaneously, i.e., without applying a flame or in 
any way inducing the reaction to take place. 

Action of chlorine on phosphorus (a non-metal). Lower a piece of 
dry yellow phosphorus into a gas-jar of chlorine. It burns spon- 
taneously, giving off white fumes of chlorides of phosphorus. 

P 4 + 6Cl 8 ->-4PCl 3 

P 4 -f 10Cl s -»-4PCl 5 


In the following cases the reactants are all metals. 

Action of chlorine on sodium. Lower a piece of burning sodium on 
a deflagrating spoon into a gas-jar of chlorine. It continues to burn, 
giving off white clouds of sodium chloride. 

2Na + CI 2 ->-2NaCl 

sodium chlorine sodium 




Action of chlorine on copper. Drop a piece of Dutch metal (a very 
thin sheet of an alloy of copper and zinc, mainly copper) into a gas- 
jar of chlorine. It burns spontaneously with a green flame to form 
copper(!I) chloride and a little zinc chloride. 

Cu + CI,-»-CuCl, 

Zn + CI, -»- ZnCl, 


Action of chlorine on iron. Place a coil of iron wire in the hard-glass 

tube in the apparatus shown in Fig. 115 and pass a stream of pure dry 

chlorine over it. On heating the wire by means of a burner, the wire 

glows and the reaction continues without application of the flame, 

black crystals of ferric chloride, or iron(III) chloride, collecting in the 

small bottle, which acts as a condenser. 

2Fe + 3Cl,-»-2FeCl 3 

iron chlorine iron(III) 


Note that the iron (III) salt is formed— an indication that chlorine 
is an oxidising agent. 

The black crystals of anhydrous iron(HI) chloride should be removed 
and placed in a desiccator, as they are very deliquescent. 

to Fume 

calcium chloride 


calcium — 

iron wire 



To prevent 
condensation of 
iron(m) chloride here 

Fio. 115. 
Preparation of iron(III) chloride. 

Chlorine as an oxidising agent 

In its oxidising behaviour, chlorine, like other oxidising agents, 
acts as an acceptor of electrons. It is converted to chlorine ions, as: 
CI f + 2«--*-2Cl- 



The electrons for this purpose are supplied by the reducing agents 
with which the chlorine is reacting. In the two cases given below, 

these are : 

(a) The ironfll) ion, which is oxidised to the ironfJIT) ion by 

electron loss, 

Fe*+ - e~ -*■ Ft?+ 

(b) The sulphite ion, which, in association with water, is oxidised 
to the sulphate ion, 

S0 3 2 " + H,0 - 2e~ -> SO t a - + 2H + 

also with electron loss. 

(a) Action on iron(IT) chloride solution. Bubble a stream of chlorine 
through a solution of iron(II) chloride (which is pale green in colour). 
The colour changes to yellow, and on adding a little caustic alkali 
solution there is obtained a reddish-brown precipitate of iron(III) 
hydroxide, showing that the iron(II) chloride has been oxidised to 

iron(III) chloride. 

2FeCl 2 -r-Cl,-»-2FeCl 3 
or 2Fe* + + Cl I ^-2Fe^-f-2Cl- 

(b) Oxidation of sulphurous acid to sulphuric acid. On bubbling 
chlorine into a solution of sulphurous acid in water for a few minutes, 
dilute sulphuric acid is obtained. 

2H,0 + SO, + CI, -^ H a S0 4 + 2HC1 
or H,0 + SO, ^ H,S0 3 ^ 2H+ + SO, 1 - 

SOs*- + H s O + Cl,-> S0 4 »- + 2H+ + 2C1- 
The presence of the sulphuric acid can be shown by testing with 
dilute hydrochloric acid and barium chloride before and after the 
experiment. There is obtained after the experiment a white precipitate 
of barium sulphate. [N.B. The sulphurous acid solution must be fresh, 
otherwise it will contain a certain amount of sulphuric acid due to 
atmospheric oxidation.] 

The displacing action of chlorine 
Chlorine can displace bromine and iodine from bromides and 

iodides. . . , . . , 

Displacement of bromine. Bubble chlorine into a solution of potas- 
sium bromide in water. The clear solution immediately turns red 
(due to formation of bromine water) and finally a drop of a red 
liquid (bromine) is observed at the bottom of the boiling-tube. 
2KBr + Cl,-v2KCl-f Br, 
Displacement of iodine. The above experiment is repeated with 
potassium iodide solution. The clear solution turns to the character- 
istic dark brown 'iodine' colour and finally a black solid (iodine) is 



deposited. On warming the solution the characteristic violet vapour 
of iodine is seen. 

2KI + Cl,-*-2KCl + I t 
In these two reactions, chlorine oxidises the ions, Br" and I - , by 
attracting the electrons from them, and chlorine is reduced to its ions 
by electron gain. 

2Br- (or 21") + Cl 8 -*■ Br 2 (or LJ + 2C1- 
Act ion of chlorine on the caustic alkalis 

On the cold dilute aqueous solution. Chlorine is absorbed by a solu- 
tion of water of sodium hydroxide or potassium hydroxide, forming 
a pale yellow solution of the hypochlorite and chloride of the metal. 
Cl 2 -f 2NaOH -* NaOCl + NaCl + H 2 
Cl 8 + 2K.OH -> KOC1 + KC1 + H 2 
or d a + 20H- ->■ OCI- + CI" + H a O 

On the hot concentrated aqueous solution. If chlorine is passed into 
a hot concentrated solution of potassium hydroxide for some time, 
there is formed a mixture of potassium chloride and potassium chlor- 
ate, and the latter can be obtained by crystallising the mixture when 
crystals of potassium chlorate separate first. (These can be purified 
by recrystallisation.) 

6KOH + 3C1 2 -> KC10 3 + 5KCI + 3H a O 
potassium potassium potassium 

hydroxide chlorate chloride 

or 60H- + 3C1 8 -> C10 3 - + 5C1 - + 3H.O 

A similar action is observed if hot concentrated sodium hydroxide 

solution or milk of lime is substituted for the potassium hydroxide 


Notice that in the above actions the alkalis are dissolved (or sus- 
pended) in water. 

Bleaching powder 

If chlorine is passed for a considerable time over solid slaked lime 
the product is bleaching powder, formerly used as an easily trans- 
ported substitute for free chlorine. Now, however, liquid chlorine is 
regularly transported in bulk and bleaching powder is less important. 
Laboratory preparation. Take a gas-jar full of chlorine and shake a 
teaspoonful of freshly prepared slaked lime into it. The colour of the 
chlorine disappears immediately. The product may be used to absorb 
the chlorine from several more gas-jars before absorption is complete. 
Ca(OH) 2 + CI, -»- CaOCl„.H 2 
calcium chlorine bleaching 

hydroxide powder 



Manufacture. Slaked lime is moved forward by Archimedean screws 
through a series of pipes against a counter-current of chlorine until 
the requisite weight of chlorine has been absorbed. The solid is 
removed and packed. 

It contains about 36% of available chlorine, i.e., chlorine which 
can be removed by dilute acids and even by the carbonic acid of the 
atmosphere. Hence bleaching powder usually smells of chlorine and 
deteriorates if in contact with air. It has an extensive use in dye works, 
laundries and even in the home. 

CaOCl 2 + CO, -* CaC0 3 + Cl 2 
bleaching carbon chlorine 

powder dioxide 

Chlorine in chemical industry 

(a) Chlorine is extensively used as a bleaching agent and in the 
manufacture of bleaching agents and domestic antiseptic solutions 
such as sodium hypochlorite. It is also used in the manufacture of 
chlorates, used, for example, as weed-killers, and in the manufacture 
of hydrogen chloride. This gas is used in the production of PVC 
(p. 461) so that chlorine is indirectly involved in this product. 

(b) Many organic chemicals are manufactured with the help of 
chlorine, e.g., carbon tetrachloride, CC1«, trichloroethylene, C 2 HCI 9 , 
chloral, CCl 3 .CHO and many others. These compounds are useful 
as degreasing agents, dry-cleaning fluids and as sources for other 
products, e.g., chloral for DDT. 

(c) Chlorine is used in sterilising water for domestic and industrial 
use and for swimming-baths. 

Chlorine occurs as an important by-product of electrolytic cells 
which produce sodium hydroxide from the raw material, sodium 
chloride. For this, see p. 253. The products appear in the mass pro- 
portion of NaOH : CI, i.e., 40 : 35.5. 

In the older Deacon's Process, hydrogen chloride (produced from 
sodium chloride and concentrated sulphuric acid) was oxidised by air 
to chlorine at about 450°C in the presence of the catalyst, copper(II) 
chloride, CuCl„ distributed on broken brick to increase the available 

4HC1 + O t -*• 2H a O + 2C1, 

This process (introduced in 1868) was apparently rendered obsolete 
by electrolytic processes about 50 years later but has been recently 

(For an account of chlorine as a member of Group VII of the 
Periodic Table, see p. 85. For chlorine molecule, see p. 375.) 
{Questions on page 385) 



Chapter 27 

Hydrogen Chloride and the Chlorides 


THE gas is usually called hydrogen chloride, or hydrochloric acid 
gas, whereas the solution of the gas in water (which when satur- 
ated contains about 36% by weight of the gas) is termed hydrochloric 


Common salt (sodium chloride) is placed in a flask fitted with a 
thistle funnel and delivery-tube, and concentrated sulphuric acid is 
added (Fig. 116). There is effervescence, and misty fumes are ob- 

Common Salt ' 
and cone. 

\ Concentrated 
Sulphuric acid 
to dry the gas 

Fio. 116. 

Preparation of hydrogen chloride. 

served. The gas is passed through a wash-bottle containing concen- 
trated sulphuric acid to dry it, and collected as shown by downward 
delivery, the gas being denser than air. 

H*S0 4 (l) + NaCI(c) -*• NaHS0 4 (c) + HCl(g) 
or H s S0 4 + CI- ->• HS0 4 - + HC1 


Fio. 117. 

Method of dissol- 
ving a very soluble 
gas in water. 

The reaction proceeds in the cold, although a further yield of the 
gas was obtained in the industrial process by heating to a red heat. 

NaCl + NaHS0 4 -> Na a S0 4 + HC1 
sodium sodium sodium hydrogen 

chloride hydrogen sulphate chloride 


(The above indicates clearly the acid nature of sodium hydrogen 

A solution of the gas in water can be made by 
means of the funnel arrangement as in Fig. 117. 
This solution is hydrochloric acid. If the gas is 
passed into water until no more gas is absorbed, 
the product is concentrated hydrochloric acid 
and contains about 36% by weight of hydrogen 

This device of passing a gas into water by 
means of an inverted funnel is essential when 
the gas is very soluble in water. If the gas is suffi- 
ciently soluble, it may be absorbed in the water 
more quickly than it is being generated in the 
flask. In this case, the pressure in the delivery- 
tube and flask is reduced and atmospheric pressure from outside 
then forces the water back up the delivery-tube. This effect is called 
'sucking-back'. If the tube is made of ordinary narrow glass 
tubing, the water will quickly fill it and pass over into the generating 
flask. This would in any case stop the reaction and might, if hot 
concentrated sulphuric acid was being used, be dangerous. When the 
inverted funnel is used, the water may begin to suck back, but a 
considerable volume of it is required to fill the funnel before the 
narrow tube is reached. This lowers the level of the water in the 
beaker, and, if the rim of the funnel was at first only just immersed, 
it will be exposed, and air will be forced under it, before the funnel 
is filled. As soon as air enters, the water drops back from the funnel 
to the beaker and the process begins again. Notice that, to be effective, 
the funnel must be arranged with its rim only just immersed. An 
additional advantage is that the funnel offers a large water surface 
for absorption of the gas. 

Test for the gas 

It is a clear gas (although in damp air it appears misty), acid 
to litmus, and produces a white precipitate of silver chloride m a 
drop of a solution of silver nitrate and nitric acid which is held on a 
glass rod in the gas. 

Ag+(aq) + Cl"(aq) -»- AgCl(c) 




It has a choking, irritating smell, and is an acid gas, which is very 
soluble in water. These latter two properties are neatly shown by the 
fountain experiment (Fig. 138). In this case, blue litmus is placed in the 
trough and turns red in the flask, showing the acidity of hydrogen 
chloride in aqueous solution. 

Hydrochloric acid 

This is a solution of hydrogen chloride in water. Being almost 
completely ionised in dilute solution, this acid is very strong, and 
shows the usual acidic properties (p. 223), as : 

HC1 -*■ H+ + CI- 

(i) liberation of hydrogen with certain metals (Mg, Zn, Fe): 

Zn + 2H+ -*■ Zn* + + H g 

(ii) neutralisation of bases to form salts and water 

Na + OH- + H + CI- -> Na+ Cl~ + H s O 

(iii) liberation of carbon dioxide from carbonates 

2H- + C0 3 a -->H s O + C0 2 

In (i), notice, however, that hydrochloric acid does not react with 

It may be mentioned here that the situation is not quite as simple 
as the above implies. Pure water-free HCI shows no acidity because 
it contains no ions. It is a covalent compound. The aqueous solution 
contains ions. The ion responsible for acidity, however, is not I he 
simple H + ion, but a hydrated form of it, H 8 0+, produced by a 
reaction between HCI and water: 

H 8 + HCl-»-HaO + + a- 
Similarly, water-free H s SO« and HNO s are probably not acidic (being 
unionised), but ionise in water and then become strongly acidic: 
H s S0 4 + 2H 2 ->- 2H a O * + S0 4 *~ 
HNO s + H a O ->- H a O+ + N0 3 - 
For ordinary purposes the hydration of the hydrogen ion is often 
ignored and H+ is used. (For HCI in toluene, see p. 240.) 

Action of oxidising agents. Chlorine is the product of oxidation of 
hydrochloric acid and most oxidising agents will liberate chlorine 
from it (see p. 355). 

Action of heat. If a concentrated hydrochloric acid solution is 
heated, hydrogen chloride escapes into the air. If a very dilute solu- 
tion is heated, water is lost, making the acid more concentrated. In 
both cases a mixture is finally obtained containing 20.24% of hydro- 



chloric acid, which distils over unchanged. This is termed a 'constant 
boiling mixture'. 

Manufacture of hydrogen chloride 

Until recently, hydrogen chloride was manufactured exclusively by 
heating common salt with concentrated sulphuric acid. This was the 
first stage of the Leblanc process. 

2NaCl + H,SO t -»• 2HC1 + Na 2 S0 4 
The gas was absorbed in water to form hydrochloric acid. 

Recently, improved methods of manufacture of hydrogen have 
made it possible to produce hydrogen chloride in quantity by direct 
combination of hydrogen with chlorine, which is obtained by electro- 
lysis of brine. 

H 2 + Cl 2 -*■ 2HC1 
This method is likely to increase in importance. 


Metallic chlorides 

Preparation. All metals are attacked by chlorine to form chlorides. 
The methods of preparation are summarised below. 

The chlorides of these metals are made by the action of: 

the alkali (or oxide) or fK 

carbonate on dilutes Ca 1 

hydrochloric acid. iNa 

(Ma 1 
the metal, oxide, or 

carbonate on dilute< 

hydrochloric acid. 

by double decom-, 

the oxide or carbon- 
ate on dilute hy-] 
drochloric acid. 

by double decom- 

Al l 
Zn 1 
Fe 1 





These metals are 
attacked by dilute 
hydrochloric acid. 

These metals not 
attacked by dilute 
hydrochloric acid. 

All attacked by 
chlorine to form 
a chloride. 

N.B. IronfJII) 1 chloride is made by the action of chlorine on the 

» The chlorides of these metals are very deliquescent. The above is a list of 
common metals in the order of the electrochemical scries. 



Aluminium chloride, A1CI 3 

This is a pale yellow solid. Being readily hydrolysed by water, it 
fumes in damp air with evolution of hydrogen chloride. If required 
anhydrous, it must be prepared by healing aluminium foil in dry 
chlorine or dry hydrogen chloride. The apparatus is the same as for 
iron(III) chloride (p. 362). 

2AJ + 3C1, -> 2A1C1 3 
2AI + 6HC1 -v 2A1C1 3 + 3H„ 

The anhydrous solid reacts rather violently with water. It forms the 
hydrate, A1C13.6HS.O, and, with excess water, dissolves and hydro- 
lyses considerably. 

AICI3 + 3H,0 ?* Al(OH) 3 + 3HC1 

On evaporation to dryness, the solution leaves hydroxide as residue. 

Ammonium chloride (sal-ammoniac), Nil .CI 

This compound has many uses, two of the chief being as a consti- 
tuent of the Leclanche" electric battery, and as a flux in soldering. It 
is usually prepared by boiling ammonium sulphate solution with 
common salt. 

(NH 4 ) 2 S0 4 + 2NaCl «* 2NH 4 C1 + Na 3 SO, 

On cooling, sodium sulphate crystallises first and may be removed by 
filtration, after which ammonium chloride may be obtained from the 
filtrate as a white solid. It sublimes on being heated. For the action 
of heat on ammonium chloride see p. 433. 

Potassium chloride, KC1 

This occurs as carnallite (K.Cl.MgCl 2 .6H 3 0). Potassium chloride 
can be prepared in the laboratory by the neutralisation of potassium 
hydroxide solution with dilute hydrochloric acid. 

KOH + HC1 -»■ KC1 + H 8 

Eotassium dilute potassium water 

ydroxide hydrochloric chloride 

It forms white cubic crystals similar to those of sodium chloride. It 
imparts a lilac colour to the non-luminous Bunsen flame. It is not 

Sodium chloride, NaCl 

This occurs as a rock salt, which is mined as solid or pumped out 
of the ground as brine. It is prepared by similar methods to those 
used for potassium chloride and forms white cubic crystals. It im- 
parts a golden yellow colour to the Bunsen flame. Pure common salt 



is not deliquescent (the dampness of ordinary salt is due to impurities, 
e.g., magnesium chloride which is deliquescent). 

Sodium chloride is a very important chemical. As a sodium com- 
pound, it is converted into caustic soda, washing soda, baking soda 
and salt-cake (sodium sulphate) and other less important sodium 
compounds. It is used to 'salt-out' soap. As a chloride, it yields 
hydrochloric acid and chlorine, used as a bleaching agent and in 
manufacture of hypochlorite solutions for home use, as well as in 
fine chemical manufacture. 

The alkali industry of south Lancashire and Cheshire is close to 
the Lancashire coalfield and to common salt deposits at Northwich 
and to limestone outcrops. The neighbouring Lancashire textile area 
takes its products to manufacture soap and for bleach and dye in- 
dustries; there is also glass manufacture in the area. The Mersey and 
its canal systems, and the Manchester ship canal, provide good bulk- 
transport facilities. 

Sodium chloride crystallises as a 
face-centred cube (see diagram). In 
an end face of the cube as shown, a 
Na + ion occupies the centre; then 
the four corners of the face are also 
occupied by Na + ions, with four Cl - 
ions spaced equally between them. 
In the next face, the positions of 
Na + and Cl~ ions are reversed, and 
so on alternately. The attractive forces between the ions are rela- 
tively great. The only ionic motion is some vibration, consequently 
the solid appears rigid and has negligible electrical conductivity. 
When heated the ions vibrate increasingly; at 801°C the lattice col- 
lapses, i.e., the salt melts and the ions become mobile. The fused salt 
conducts electricity by migration of Na + to cathode and Cl~ to 
anode where each is discharged to give sodium and chlorine. 

2Na + + 2e~ -*■ 2Na; 2C1" -»■ Cl 2 + 2e~ 
The lattice arrangement in the solid is determined by X-ray diffraction. 
A narrow beam of X-rays is passed through a crystal of the com- 
pound and on to a photographic plate. After development, the plate 
shows a central spot, produced by X-rays which have passed straight 
through the crystal, and, round it, a diffraction pattern of spots from 
which the lattice arrangement can be deduced. 

Calcium chloride, CaCL 

This is very deliquescent and can be used as a drying agent for 
most gases (but not for ammonia, with which it forms a compound). 

o -ci~ 

Crystal lattice of Sodium Chloride 



The anhydrous salt is prepared by evaporating a solution until the 
solid formed fuses. The solution is most easily prepared by adding 
marble or limestone to dilute hydrochloric acid until a little of the 
marble remains. The mixture is then filtered. 

CaCO a + 2HC1 -► CaCl, + H g O + CO, 

Anhydrous ferric chloride, iron(lll) chloride, FeCl 3 

This is made by the action of iron on chlorine by the method 
described on p. 362. 

The anhydrous salt cannot be made by the evaporation of the 
solution because the chloride is attacked by water when a concen- 
trated solution is evaporated. This type of action is termed hydrolysis. 

/C\ HOH 
Fef-jCl + HOH 

\jCl HOH \OH 


+ 3HClf 

loses wa ter on heating 

Fe,0 8 

On heating the solution in air the final product is iron(III) oxide. 

This action is the reverse of neutralisation. 

Iron(III) chloride is a black solid in the anhydrous state, but forms 
a brown solution if concentrated, and a yellow solution if dilute. It 
can be reduced by reducing agents [e.g., zinc and dilute hydrochloric 
acid) to iron(II) chloride. 

Zn(c) + 2Fe a+ (aq) -> Zn 2+ (aq) + 2Fe 8+ (aq) 

With alkaline solutions, it gives (as do all iron(ITI) salts dissolved in 
water) a reddish-brown gelatinous precipitate of iron(III) hydroxide. 

FeCl a + 3NaOH -*■ Fe(OH) 3 + 3NaCl 

ionically: Fe 8+ (aq) + 30H-(aq) -> Fe(OH) 8 (c) 

Ferrous chloride, iron(IT) chloride, FoCl, 

Anhydrous iron(H) chloride, a white solid, is made by heating iron 
wire strongly in a stream of dry hydrogen chloride. 

Fe + 2HCl-»-FeCl, + H t 
iron hydrogen iron(II) hydrogen 
chloride chloride 

It forms a pale green solution which gives, with alkaline solutions a 
dirty green precipitate of iron(II) hydroxide (as will any iron(U) salt 
dissolved in water). 



FeCl, + 2NaOH 



2NaCl + Fe(OH), 

ionically: Fe 2+ (aq) + 20H"(aq) -> Fe(OH) 2 (c) 

Lead(ll) chloride, PbCl 8 

This is a white insoluble substance made by the interaction of a 
solution of any soluble lead(II) salt with a solution of any soluble 

Into a beaker put some dilute hydrochloric acid and add lead(U) 
nitrate solution. There is a white precipitate. 

PrXNOs)* + 2HC1 -»- PbCl a + 2HNO, 

ionically: Pb 2+ (aq) + 2Cl-(aq) -*• PbCl 2 (c) 

Filter off the white precipitate, wash it two or three times with a little 
cold distilled water and put it on a porous plate to dry. 

Lead(II) chloride is almost insoluble in cold water yet fairly soluble 
in hot water. (It is the only common substance which shows this 
peculiar behaviour.) 

Silver chloride, AgCl 

This is a white insoluble compound made by adding a solution of 
any soluble silver salt to a solution of any soluble chloride. 

AgNOg + NaCl - 
silver sodium 
nitrate chloride 

► AgCll + NaN0 3 

silver sodium 

chloride nitrate 

ionically: Ag" 1 (aq) + Cl-(aq) -> AgCl(c) 

The white solid is filtered off, washed two or three times with hot 
distilled water and dried on a porous plate. (The whole action should 
be performed in the absence of light since silver chloride turns violet 
on exposure to light.) 

All the silver halides, AgCl, AgBr and Agl, emit electrons when 
exposed to light (a photoelectric effect) and are slowly reduced, turn- 
ing violet and, eventually, black. The chief reduction product is 
metallic silver. 

Ag+ + e~ — *■ Ag 

The halides show decreasing photoelectric activity in the order: 
AgBr -* AgCl — *■ Agl. This is why silver bromide is used as the 
principal silver compound suspended in the gelatine of a photo- 
graphic plate. The other silver halides are less sensitive to the action 
of light. 



Properties of chlorides 

(1) Action with concentrated sulphuric acid. On being treated with 
concentrated sulphuric acid, a chloride evolves hydrogen chloride, 

NaCl(c) + H 2 S0 4 (1)-* NaHSO«(c) + HCl(g) 

(2) Chlorides are more volatile than most salts. 

The chlorides are on the whole a volatile class of compounds. This 
makes them suitable for use in the 'flame-test' in which certain metals 
can be detected by the colour their vapour imparts to the Bunsen 
flame. To perform the flame-test, the substance under consideration 
is moistened with concentrated hydrochloric acid and a nichrome 
or platinum wire is dipped into the mixture and applied to the non- 
luminous Bunsen flame. 

Metal chloride 




Persistent golden yellow 

(invisible through blue glass) 

Lilac flame 
(visible through blue glass) 


Green (blue zone) 


(3) Hydrolysis of chlorides. Several chlorides are readily hydro- 
lysed by water, e.g., magnesium, zinc and iron chlorides. If solutions 
of the chlorides are evaporated a basic salt or the oxide of the metal 

(4) Action of concentrated sulphuric acid on mixture of chloride and 
oxidising agent. Mix together a little common salt and manganese 
dioxide (many other oxidising agents would be suitable). Put this 
into a test-tube, add a few drops of concentrated sulphuric acid, and 
warm. A green gas, chlorine, is evolved. 

2NaCl + 2H 2 SO« + MnO a -»■ 
sodium concentrated manganese 
chloride sulphuric dioxide 

acid manganeseflV) 


MnSO« + Na,SO« 

manganese sodium 

sulphate sulphate 

+ 2H.O + a, 

water chlorine 

Test for a soluble chloride 

Dissolve a suspected chloride in distilled water and add a little 
nitric acid and then silver nitrate solution. If a chloride is present you 
will see a white precipitate of silver chloride. 

Ag + (aq) + Cl-(aq) -> AgCI(c) 



Divide the precipitate into two parts. Add ammonium hydroxide to 
one and observe that the precipitate dissolves. Allow the other to be 
exposed to the light for a few minutes. The precipitate will turn violet. 

Silver chloride is insoluble in nitric acid but soluble in ammonium 

The only two common insoluble chlorides arc lead chloride and 
silver chloride. 

The chlorine molecule 

It is known by experiment (p. 112) that 1 volume of chlorine 
combines with hydrogen (also 1 volume) to produce 2 volumes of 
hydrogen chloride in constant conditions. Applying Avogadro's Law 
to the chlorine and hydrogen chloride, 1 molecule of chlorine pro- 
duces 2 molecules of hydrogen chloride. Since the least possible 
amount of chlorine in a molecule of hydrogen chloride is one atom 
(atoms being indivisible for chemical purposes), the molecule of 
chlorine must supply at least two atoms to the hydrogen chloride. 
That is, the chlorine molecule is at least diatomic. It is also known 
that chlorine has never produced more than twice its own volume 
of any gaseous chlorine compound in constant conditions. This 
points to an actually diatomic molecule for chlorine, i.e., C\ t . 

Further, the specific heat of gaseous chlorine at constant pressure 
is greater than its specific heat at constant volume in the ratio of 
1.36: 1. This is close to the recognised value of 1.40 for diatomic 

Questions on this chapter will be found on page 385. 

Chapter 28 

Bromine and Iodine 


BROMINE was discovered by Balard in 1826. He passed chlorine 
through the mother liquor obtained after crystallising common 
salt from sea-water. The liquor turned red and from it he was able to 
isolate bromine and to show that it was an element. 

Liebig, some years previously, had received the dark red liquid with 
a request to examine it, but thinking that it was merely a compound 
of iodine and chlorine, he did not pay it much attention. 


Bromine occurs chiefly as the bromides of potassium, sodium and 
magnesium, usually in association with larger proportions of the 
chlorides of those metals. Since the bromides are much more soluble 
in water than the chlorides, a liquid rich in bromides is left by 
crystallising out the chlorides. Treated in this way, the mother liquors 
from the Stassfurt deposits in Germany (which consist mainly of 
carnallitc, KCl.MgCl2.6HjO), after the removal of a large proportion 
of the potassium chloride contain about |% of bromide and the 
bromine is obtained by allowing this solution to come into contact 
with chlorine, which displaces the bromine: 

MgBr. + Cl,- 

MgCl 2 + Br, 

Laboratory preparation of bromine 

This experiment must be done in a fume-chamber. 

Bromine can be prepared in a way exactly analogous to one of the 
methods for making chlorine. 

Make an intimate mixture of potassium bromide and manganese 
dioxide (manganese(IV) oxide) and place this in a retort (Fig. 1 18). 
Add some concentrated sulphuric acid and warm the mixture. A red 
gas is given ofT(together with some misty fumes of hydrogen bromide) 




which condenses to a red liquid in the cooled receiver. This is 

2KBr + MnO, + 2H a S0 4 -► K,S0 4 + MnSO« + 2H s O + Br, 
potassium manganese concentrated bromine 

bromide (IV) oxide sulphuric 

Cone Sulphuric add. 
•Manganese Dioxide, 
and Potassium Bromide. 

Fio. 118. 
Preparation of bromine. 

Physical properties of bromine 

(1) It is a heavy (density = 3.2 g/cm 3 ), red, volatile liquid (boiling- 
point 59°C). 

(2) It has a choking, irritating smell. (Its name means 'a stench'.) 
The liquid causes sores on the flesh, which heal with difficulty. 

(3) It is slightly soluble in water, forming a red solution containing 
about 3% of bromine at ordinary temperatures. 

The following chemical properties are considered under the same 
headings as those of chlorine. 

Bromine as a bleaching agent 

Bromine is a bleaching agent, not so rapid as chlorine. A piece of 
damp litmus paper is bleached when placed in the vapour of bromine. 

Bromine with hydrogen 

Bromine combines with hydrogen, but not as readily as does chlor- 
ine. A mixture of chlorine and hydrogen will explode when merely 
exposed to sunlight, but a mixture of bromine and hydrogen needs 



the application of heat to induce combination, and the compound 
formed (hydrogen bromide) is not as stable as hydrogen chloride. 

H, + Br,->-2HBr 

Bromine as a bromide former 

Bromine combines readily with most metals and non-metals to 
form bromides, for example, copper, iron, sodium, sulphur. It ex- 
plodes when mixed with yellow phosphorus, so vigorous is the action. 
Phosphorus tribromide is made by gradually adding a solution of 
bromine in carbon tetrachloride to red phosphorus. The solution is 
used in order to moderate the action. 

Bromine as an oxidising agent 

Bromine is an oxidising agent, Br, -f 2e~ — *■ 2Br _ , but not quite 
as vigorous as chlorine. It will perform the majority of the oxidations 
attributed to chlorine. Thus, on shaking an acidified solution of 
iron(II) sulphate with a few drops of bromine in a test-tube, the 
bromine colour soon disappears and the iron(II) sulphate has been 
converted into iron(IH) sulphate. 

2Fe s+ + Br, -»- 2Fe 8+ + 2Br~ 

The displacing action of bromine 

Bromine can displace iodine from iodides but cannot displace 
chlorine from chlorides. Thus, on adding a few drops of bromine to a 
solution of potassium iodide in water, the characteristic brown colour 
of the solution of iodine in potassium iodide is seen. On boiling the 
solution the violet vapour of iodine may be observed. 

2KI + Br 2 -»-2KBr + I, 
or 2I- + Br,-»-2Br- + I 1 

Action of bromine on the alkalis 

The action of bromine on an alkaline solution is exactly analogous 
to that of chlorine. Thus: 

Cold potassium hydroxide solution 
2KOH + Br,-*KBr 

+ KBrO + 
potassium potassium 
bromide hypo- 


Hot potassium hydroxide solution 

6KOH + 3Br, -»- 5KBr + KBrO, + 3H t O 





Bromine is an element very similar to chlorine but differing from it 
principally in that bromine is less active than chlorine. 


As indicated above, hydrogen bromide is not as stable as hydrogen 
chloride and is decomposed to some extent by the action of heat. 
This, together with the fact that hot concentrated sulphuric acid is 
an oxidising agent, makes it impossible to prepare pure hydrogen 
bromide by the action of heat on a mixture of concentrated sulphuric 
acid and potassium bromide. The products of such an attempt would 
be hydrogen bromide, bromine and sulphur dioxide. 

It may be made by the action of bromine on a mixture of red 
phosphorus and water. The chemistry of the action is that bromides 
of phosphorus are formed which are decomposed by the water. 

Preparation of hydrogen bromide 

Make a paste of red phosphorus and water and place this in the 
flask (sand may be added to 'dilute' the mixture). Bromine is dropped 
in gradually from a tap-funnel and the reaction proceeds at ordinary 
temperature (Fig. 119). Heat is evolved during the process and a 
considerable amount of bromine may be volatilised and would, if not 
removed, contaminate the hydrogen bromide. The bromine is 
removed by passing it through a U-tube containing beads smeared 

■ Bromine 

Stiff paste. 

of /fid 
Sard and 



Fio. 119. 

Preparation of hydrogen bromide. 



with red phosphorus and water. This U-tube is, in fact, a secondary 
generating apparatus so arranged as to offer a large area of phos- 
phorus to the bromine so that it is as completely removed as possible. 
The misty gas, very similar in appearance to hydrogen chloride, is 
collected by displacement of air as shown in Fig. 119, the gas being 
denser than air. It can also be dissolved in water to form hydro- 
bromic acid by the apparatus shown on p. 379. 

4P + 6Br, — > 4PBr 8 
phosphorus bromine phosphorus 


PBr s + 3H a O -> H 3 PO a -f 3HBr 
phosphorous hydrogen 
acid bromide 

(Alternatively hydrogen bromide may be prepared by heating 
potassium bromide with cone, sulphuric acid diluted with half its own 
volume of water.) 

Preparation of hydrogen bromide from its elements 

The gas can also be made by bubbling hydrogen through a wash- 
bottle containing bromine and passing the gases through a heated 
tube containing platinised asbestos, which acts as a catalyst. Any 
unattacked bromine is absorbed by red phosphorus. 

H, + Br 2 -*-2HBr 

Preparation of hydrogen bromide by the action of hydrogen sulphide 
on bromine-water 

A convenient method for making a solution of hydrogen bromide 
is to bubble hydrogen sulphide through bromine-water for some time. 
Sulphur precipitates and can be filtered off. 

H,S + Br 2 -)-2HBr-|-S| 

Test. Hydrogen bromide turns damp blue litmus paper 

red, and gives a pale yellow precipitate of silver bromide 

with a mixture of silver nitrate solution and nitric acid. 

The precipitate is only slightly soluble in aqueous ammonia. 

Ag'(aq) -(- Br-(aq) -* AgBr(c) 

Properties of hydrogen bromide 

(i) It is a dense fuming gas with a choking smell (density = 2.8: 
air = 1). 

(ii) It is very soluble in water, forming a strongly acid solution. A 
saturated solution of hydrogen bromide contains about 70% by 
weight of hydrogen bromide at ordinary temperatures. 



(iii) It is less stable than hydrogen chloride, being more easily 
decomposed into its elements. 

(iv) It is, in its general chemical properties, similar to hydrogen 


The bromides are prepared, generally speaking, by the same 
methods as the chlorides and possess similar properties. They can 
readily be distinguished from the chlorides by the action of chlorine 
gas which has no effect on the chlorides but displaces bromine from 
bromides (see p. 378 for experimental details). 


Iodine was discovered in 1812 by Courtois. He treated with con- 
centrated sulphuric acid the mother liquors obtained after extracting 
sodium carbonate from the ash obtained by burning seaweed (kelp). 
The ash contains a small percentage of iodides and the concentrated 
sulphuric acid formed hydrogen iodide and oxidised it to iodine. 
Gay-Lussac and Davy investigated the properties of the black solid, 
which was called iodine by Gay-Lussac. 

Most of the iodine used today occurs as calcium iodate, Ca(IO s )t, 
in the sodium nitrate deposits in Chile. The amount is very small 
(about 0.1% but after the removal of the sodium nitrate by crystal- 
lisation the proportion is much higher in the residues. The iodine is 
obtained by treatment with sodium hydrogen sulphite. 

Laboratory preparation from potassium iodide 

Grind together some potassium iodide and manganese dioxide 
(manganese(IV) oxide) in a mortar and place the mixture in a dish. 

Black crystals 
oF Iodine 

Cone. Sulphuric acid. 
Manganese Dioxide, 
and Tolassium Iodide 

Fio. 120. 
Preparation of iodine 



Add concentrated sulphuric acid and place an inverted funnel over 
the dish as shown in Fig. 120. Warm the mixture carefully and the 
violet vapour of iodine will be seen to condense on the cooler parts of 
the funnel to black shining plates. The chemistry of the action is 
similar to the formation of chlorine from common salt. The hydrogen 
iodide is, however, much more easily oxidised than even hydrogen 
bromide. (For iodine molecule and crystals, see p. 75.) 

Mn0 2 + 2K1 + 2H 2 S0 4 -»- K 2 SO« + MnS0 4 + 2H t O + I, 
potassium iodine 


Properties of iodine 

(i) It is a black shining solid. Density 4.9 g/cm -3 . 

(ii) It sublimes when heated rapidly, forming a violet vapour from 
which the black solid can again be obtained by cooling. 

(iii) It is almost insoluble in water but readily soluble in potassium 
iodide solution. This is due to the formation of a compound of 
potassium iodide and iodine, KI 3 , which readily dissolves. This 
solution is brown. It also dissolves in alcohol and ether, forming 
brown solutions, and in carbon disulphide and chloroform, forming 
violet solutions. 

The following chemical properties of iodine are considered in the 
same order as those of chlorine and bromine. 

Iodine does not bleach, and has little affinity for hydrogen. The 
effect of heating hydrogen iodide is to decompose the compound into 
its elements. 

Iodine as an iodide former. Iodine is a fairly active element and will 
combine with many metals to form iodides, but it does so much less 
readily than either chlorine or bromine. 

Oxidising action of iodine. Iodine is a mild oxidising agent. It will 
not perform many of the ordinary oxidising actions attributed to 
chlorine and bromine. It will, however, oxidise hydrogen sulphide to 
form hydrogen iodide and liberate sulphur. 

H 2 S + I 2 -*2HI + S|, 

Action of iodine on alkaline solutions. The action of iodine with 
alkalis is similar to the reactions of chlorine and bromine with 
alkalis. Hypoiodites, iodides and iodates are produced (p. 364). 

Displacing action of iodine. Iodine cannot displace chlorine or 
bromine from chlorides or bromides. 

Action of iodine with starch solution 

Place a 400 cm 3 beaker full of water on a tripod and gauze and heat 
to boiling. Make a paste of a small amount of starch (about 1 g) 



and a little water, and pour this into the boiling water and stir. Allow 
to cool, or if the starch paste is required immediately, pour some of 
the paste into a boiling-tube and cool under the tap. 

Add the smallest possible quantity of a solution of iodine (the test 
is sensitive to one part in one million) and immediately you will 
observe a blue coloration. Warm the mixture and the blue colour 
will disappear, but will return on cooling. 

This test is given only by free iodine and is not given by, say, a 
solution of potassium iodide in water. 

Uses of iodine 

The antiseptic properties of iodine have caused a large increase in 
the demand for iodine during the last few years. It is sold as 'tincture 
of iodine' — a solution of iodine in dilute alcohol. It is used as iodine 
and iodides in medicine to treat cases of goitre, which disease is 
thought to be due to lack of iodine in the body. Small amounts of 
iodine have, in fact, been shown to be essential to the human body 
and all other forms of vertebrate life. Remarkable results have been 
obtained by giving poultry a small but regular dose of potassium 
iodide. The production of eggs was wonderfully increased. 


This gas is much less stable than even hydrogen bromide. A 
solution of hydrogen iodide in water quickly darkens because of the 
formation of iodine. It is usually prepared by the action of water on a 
mixture of red phosphorus and iodine. Since iodine is a solid and 
does not volatilise to an appreciable extent during the reaction, there 
is no need for a U-tube as in the case of the similar preparation of 
hydrogen bromide (Fig. 119). 

If the tube, is heated hero 
violet vapours of 
Iodine are observed 

With Ammonh 
.forms Wi 
fumes of 

lodin*. red 
and 'a little 

Solution of 
Silver rubrate 
gives yellow 

Fig. 121. 
Simple apparatus to show some of the properties of hydrogen iodide. 



For class purposes, the reactions of hydrogen iodide can easily be 
shown by the very simple apparatus shown in Fig. 121. Grind a little 
red phosphorus with iodine in a mortar and introduce the mixture 
into a dry boiling-tube. Add about four drops of water, insert a cork 
fitted with delivery-tube, and allow the gas to fall (hydrogen iodide is 
four times as dense as air) into test-tubes, one containing silver 
nitrate solution and the other containing a few drops of ammonia. 
The tube may also be heated when the decomposition of the gas is 
obvious from the violet vapour of iodine which is observed. 

4P + 6I 2 -»-4Pl3 
PI 3 + 3H s O-»- 3HI + H 3 P0 3 
phosphorus hydrogen phosphorous 

tri-iodide iodide acid 

A solution of hydriodic acid can be more simply obtained by 
bubbling hydrogen sulphide into a suspension of iodine in water. 
The end of the reaction is reached when all the iodine is seen to have 
disappeared. The precipitated sulphur is filtered off. 

H,S + I a -*2HI-r-S| 

Test. Add a little chlorine-water to a gas-jar of the gas, 
and pour a few drops of the liquid into starch paste. A 
blue colour is observed. 

Properties. Hydrogen iodide is a fuming, choking gas, very soluble 
in water, forming hydriodic acid. The gas readily dissociates (rever- 
sibly) above !80°C. 

2HI ?*H, + l s 

The acid, usually used as acidified potassium iodide solution, is a 
vigorous reducing agent (21" — 2e~ — > T 2 ). It is oxidised with liber- 
ation of iodine (brown) by exposure to air and by : 
hydrogen peroxide 

H a 2 -f 21- + 2H+ -> 2H 2 + I 2 

potassium permanganate 

2MnO«- + 101- + 16H' ->• 2Mn 2+ + 8H a O + 51, 


Iodides are similar to chlorides and bromides but can be readily 
distinguished by the action of chlorine or bromine, which liberate 

2KI + Cl 2 ->-2KCl + I ! 

Addition of silver nitrate in dilute nitric acid to an iodide solution 

precipitates silver iodide, a yellow salt which is insoluble in ammonia. 

Nal + AgNQ 3 -*- Agl -f NaNO, 



Comparison of halogen elements, Cl„ Br, and I, (see also p. 85) 

These elements show very marked similarity of properties because all 
have seven electrons in the outermost electron shell. 

CI 2,8,7 Br 2,8,18,7 12,8,18,18,7 

They are all univalent, forming an ion by gain of one electron per atom. 
This completes the external octet. 

X, + 2e~ — > 2X", where X is CI, Br or I 

Being electron acceptors, they are all non-metals and oxidising agents; 
chlorine is the most powerful oxidising agent and iodine the least. Typical 
oxidations are: 

2Fe'+ + CI, — *■ 2Fe*+ + 2C1" (iron(II) ion to iron(III) ion) 
S»- + Br, — > S + 2Br- (H,S to sulphur) 
Because of their very marked oxidising action, chlorine and bromine are 
bleaching agents; iodine is not. 

In the order: CI — *■ Br — *■ I, each halogen displaces an element to the 
right of it from simple salts. To do this, the more powerful oxidising agent 
oxidises the ion of the other halogen, so liberating the element. 
CI, + 2Br- -*■ 2C1- + Br, 
Br, + 2I-->2Br- + I, 

All three halogens behave in a similar way with aqueous caustic alkali 

Cold dilute: X, + 20H" —*■ X" + XO~ + H.O 
Hot cone: 3X, + 6OH-— »• 5X" + XO," + 3H a O, 

where X is CI, Br or I. For names of products, see p. 364. 

The silver salts of the three halogens, AgCl, AgBr, Agl, are all insoluble 
in water and in dilute nitric acid. In the order given, they show gradation of 
colour: white — ►pale yellow —*■ yellow, and gradation of solubility in 
ammonia : very soluble — ► slightly soluble — *■ insoluble. All the silver salts 
are blackened by exposure to light, with reduction to metallic silver. 

All these halogens combine directly with hydrogen (chlorine most readily, 
iodine least) and the halides, HC1, HBr, HI, are gaseous in ordinary con- 
ditions, very soluble in water and strongly acidic. The gases show a grada- 
tion of stability, being stable up to 1500°C, 800°C and 180°C respectively. 

Chlorine exhibits a covalency of seven, i.e., the number of electrons in 
the outer shell, forming the acidic oxide, C1,0,. Neither of the other halo- 
gens shows this maximum valency towards oxygen; iodine forms the acidic 
oxide I,0„ but bromine forms no stable oxide. 


1. 'Chlorine is the product of the oxidation of hydrogen chloride.* 
Illustrate this statement by describing two different processes for the pre- 
paration of chlorine, one used in the laboratory and the other a large-scale 
process. State the various stages in the bleaching of coloured cloth by 
means of bleaching powder. (N.U.J.B.) 

2. Describe a laboratory method for preparing chlorine. What im- 
purities would the gas so prepared be liable to contain, and how would 
you get rid of them? (O. and C.) 



3. How would you prepare chlorine from bleaching powder ? What takes 
place when (a) a mixture of chlorine and hydrogen is exposed to diffused 
daylight? (6) chlorine is passed into cold, dilute potassium hydroxide 
solution ? (c) chlorine is passed into hot, concentrated potassium hydroxide 
solution? (d) chlorine water bleaches litmus solution? (D.) 

4. How can it be shown that hydrogen chloride contains half its volume 
of hydrogen? 

Contrast the behaviour of hydrogen chloride with hydrogen iodide (a) 
when heated; (6) when mixed with bromine vapour (O. and C.) 

5. Describe how you would prepare a specimen of bleaching powder in 
the laboratory. What happens when it is treated with (a) dilute hydrochloric 
acid; (b) an acidified solution of iron(ll) sulphate? Describe how you 
would employ your specimen of bleaching powder to bleach a piece of red 
cloth. (O.) 

6. Describe the preparation from iron of (a) ironfl!) chloride; 
(b) iron(III) chloride. How may these compounds be converted one into 
the other? (O. and C.) 

7. Give an account of the experiments you would do in order to in- 
vestigate the action between lead dioxide and hydrochloric acid. By what 
tests would you identify the more important products? (L.) 

8. Chlorine will react with the substances slaked lime, iron, potassium 
hydroxide, hydrogen sulphide, potassium iodide and water. 

Describe briefly the apparatus you would use in carrying out these re- 
actions and state clearly the conditions necessary to bring about the 
chemical change. Where possible, give an equation for the reaction. (D.) 

9. Describe and explain the action of chlorine on (a) metallic sodium; 
(b) sodium hydroxide solution; (c) potassium iodide solution. 

An excess of chlorine water was added to 50 cm 3 of a solution of sulphur 
dioxide. An excess of barium chloride (BaCI.) solution was then added 
and the resulting precipitate after filtering, washing and drying, weighed 
0.5202 g. Calculate the volume of sulphur dioxide at s.t.p. which was 
dissolved to make the original solution. (Ba = 137; S = 32.) (L.) 

10. Suggest simple experiments by which you could show that (a) 
ammonia contains hydrogen; (_b) carbon dioxide contains carbon; and (c) 
common salt contains chlorine. (N.U.J.B.) 

11. (a) Describe with a sketch how you would prepare and collect some 
dry chlorine. 

(b) Assuming that chlorine can be completely converted into iron(III) 
chloride (FeCl a ) by passing the gas over heated iron, calculate what volume 
of chlorine, measured at I7°C and 870 mm pressure, would be required to 
produce 32.5 g of iron(IH) chloride. (Fe = 56, CI = 35.5, 2 g of hydrogen 
occupy 22.4 dm 8 at 0°C and 760 mm pressure.) (N.U.J.B.) 

12. How would you obtain from common salt a concentrated solution 
of hydrogen chloride? 

Explain, giving equations, what takes place when the solution is treated 
with (a) ammonia; (b) manganese dioxide; (c) iron(II) sulphide. (O. and C.) 

13. (a) Indicate the chemistry of on b process of manufacture of each of 
the following: (i) calcium oxide; (ii) chlorine; (iii) bleaching powder. (A) 
What takes place when carbon dioxide is passed into a suspension of 
bleaching powder? (N.U.J.B.) 

14. Describe how you would prepare and collect a small quantity of 



liquid bromine. Under what conditions does this element combine with 
hydrogen? and how does the compound so formed resemble and differ 
from the corresponding compounds of chlorine and iodine I (L.) 

15 How would you make an aqueous solution of hydrochloric acid 
starting from sodium chloride? Sketch the apparatus you would use What 
tests would you apply to show that the solution you have made is (a) aaa% 
(5 hydrochloric acid ? Give one test in each case Hot -would you determine 
the volume composition of hydrogen chloride ? (C.W.B.) 

1 6 Describe the preparation and the collection of dry chlorine. Explain 
the bleaching action of chlorine and contrast it with that of sulphur 
How would you remove an ordinary ink stain from a white cotton cloth by 
means of bleaching powder? (L.) 

17 Explain in electronic terms the displacement of bromine, by 
chlorine, from aqueous potassium bromide solution. Why can i this 
reaction be regarded as a redox (oxidation-reduction) reaction? Why, in 
Sonic ternTs, does iodine not displace, chlorine from chlondes? 
Astatine, At, being in Group VII of the periodic table and in a later period 
Than iodine, what reaction would you expect between iodine and a 
metallic astatide? Justify your expectation. 

18. Quote the arrangement of electrons in the atoms of chlorine, 
bromine and iodine. Explain in electronic terms the formation of (a) the 
ion Cl ! , (Sthe molecule, I„ (c) the molecule POC1, W) the molecule 
CC ,. Briefly compare the oxidising power of these three halogen atoms 
as shown in their behaviour with hydrogen. State wi h reasons what you 
would expect as the behaviour of hydrogen with the element astatine, 
which has^tomic number of 85 and is in Group VII of the periodic table. 

Chapter 29 


(For sulphur in periodicity, see p. 87) 

THE element, sulphur, is a yellow solid. It is usually sold as either 
•flowers of sulphur', a powder, or 'roll sulphur', cylindrical sticks. 

Uses of sulphur 

The output of sulphur in the world today exceeds 32 million tonnes 
annually, about two-thirds of it being produced by the United States 
of America. This vast amount is used in the following ways: 

1. For the manufacture of sulphuric acid (see p. 411). 

2. For dusting vines to prevent the growth of certain kinds of 

3. In making calcium hydrogen sulphite, Ca(HSO s )„ which is 
used as a bleacher of wood-pulp in the manufacture of paper. 

4. For the vulcanisation of rubber, a process which converts the 
soft pliable rubber into the hard, tough substance of which 
motor tyres and similar products are made. 

Fio. 122. 

American sulphur deposits. 




5. In smaller quantities for the manufacture of dyes, fireworks, 
sulphur compounds, such as carbon disulphide, CS t , and 
medicinally in ointments. 


Sulphur occurs: 

1. In Louisiana and Texas, U.S.A., as free sulphur; 

2. In petroleum gases, e.g., at Lacq, as hydrogen sulphide. 

Extraction of sulphur 

In America the deposits lie at a depth of about 160 m with 
deposits of limestone, clay and sand between the ground level and 

' <— Not compressed air 

, [Molten Sulphur 

r r=-*-\ana' Water 

'Super-heated voter 
170'C under 

rater*-, j 



and n 









Fio. 123. 
Frasch sulphur pump. 

the sulphur (Fig. 122). It is not necessary to mine the sulphur by 
sinking shafts as in the case of coal, for sulphur differs from coal in 
having a fairly low melting-point (115°C). By utilising this property 
the sulphur can be extracted, by a method invented by Frasch, 
cheaply, rapidly and in a high state of purity. 

A hole about 30 cm in diameter is bored down through the clay, 
sand and limestone to the sulphur beds. This boring is lined with an 



iron pipe and, inside the pipe, is sunk a device called the sulphur 
pump. It consists of three concentric tubes which terminate in a 
reservoir of larger diameter (see Fig. 123). Down the outermost of the 
three tubes is forced a stream of water at about 170°C. This water 
must be kept at a pressure of about 10 atm per square inch to main- 
tain it in the liquid state, i.e., it is super-heated water, and it is hot 
enough to melt the sulphur. The molten sulphur flows into the 
reservoir at the base of the pump and is forced up to the surface 
through the second of the three tubes by means of a blast of hot 
compressed air at a pressure of about 15 atm., which is forced down 
the narrowest tube. The sulphur is run into large tanks, where it 
solidifies and can be separated from the water. Sulphur more than 
99% pure is produced by this operation and a single pump may 
produce up to 500 tons of this high-grade sulphur daily. 

Sulphur from petroleum gases. At Lacq (S. France), gas associated 
with petroleum deposits contains mainly methane, CH«, with carbon 
dioxide and hydrogen sulphide (15%). It is passed, at 70 atm pres- 
sure, over an (alkaline) amine solution, which absorbs the acidic 
carbon dioxide and hydrogen sulphide. These gases are then released 
by heating at atmospheric pressure. Controlled supplies of air are 
added and the hydrogen sulphide is oxidised to sulphur in three 
stages, the last two in contact with a heated catalyst, bauxite. 

2H a S + 0,->2H t O + 2S 
The sulphur vapour is condensed and then cooled to solid. The 
product is 99.9% pure. The deposits have been worked since 1957 
and present yield is about 1.4 million tonnes per annum. 

The action of heat on sulphur when air is excluded 

Place some powdered roll sulphur in a narrow test-tube and 
warm it gently, shaking well. Try to avoid local over-heating by rotat- 
ing the test-tube. The sulphur passes through the following stages as 
the temperature rises: 

1. It melts at about 115°C to an amber-coloured, mobile 

2. It becomes much darker in colour and, suddenly, at 160°C, 
very viscous. So viscous does it become that the test-tube may 
be inverted without loss of sulphur. 

3. The sulphur gradually becomes more mobile again and very 
dark reddish brown in colour. 

4. The sulphur boils at 444°C, giving off light brown sulphur 

These changes occur in the reverse order as the sulphur cools. 



The action of heat on sulphur with a plentiful supply of air 

Plunge a deflagrating spoon containing burning sulphur into a gas- 
jar of air. The sulphur burns with a blue flame and leaves a misty 1 

Treat several gas-jars in this way and use them for the following 


Add blue litmus solution. 

It is turned red. The gas is an acidic oxide. 
Add a dilute (pink) solution of potassium permanganate. 

It is turned colourless. 
Add a dilute (golden yellow) solution of potassium dichromate. 

It is turned green. 

The results of these tests prove that the gas is sulphur dioxide (see 
p. 403). Sulphur burns in air, forming sulphur dioxide. 

S + 0,->-SO a 

Formation of sulphides from sulphur 

Sulphur will combine directly with many elements forming sul- 
phides. For example, if a finely ground mixture of iron filings and 
sulphur, in the proportions of 56 to 32 by weight (Fe = 56; S = 32), 
is heated, the two elements will combine vigorously and the whole 
mass will glow spontaneously when once the combination has been 
started at one point. A black, or dark grey, residue of iron(lI) 
sulphide is left. 

Fe + S -v FeS 

Hot copper foil or wire will similarly glow in sulphur vapour, forming 
copper(I) sulphide, Cu.S. 

2Cu + S — >■ Cu,S 
Carbon combines directly with sulphur to form the important liquid, 
carbon disulphide, CS,. 

C + 2S->CS, 

A very high temperature is required to bring about the combination, and 
this is secured by means of the electric furnace, in which an electric arc is 
struck between carbon electrodes and raises coke to white heat. Sulphur is 
also fed into the furnace. It vaporises and combines with the white-hot 
coke. Carbon disulphide vapour passes off and is condensed. 

Carbon disulphide is poisonous and may be used to destroy low and 
harmful forms of life, such as grain weevils (which feed on stored grain) or 
cockroaches. It is very inflammable, and must be used with care. It is also 
an excellent solvent. (See also CC1,, p. 297.) 

1 The misty effect is due to traces of sulphur trioxide formed simultaneously. 






Action of adds on sulphur 

Dilute acids do not act upon sulphur. It is oxidised by hot con- 
centrated sulphuric acid with formation of sulphur dioxide. 
S + 2H,SO, -»- 3SO» + 2H s O 

In this reaction the sulphur is oxidised by the acid to sulphur 
dioxide and the acid is reduced to the same substance. Of the three 
molecules of sulphur dioxide in the equation, one is the product of 
oxidation of sulphur and two are the products of reduction of the 
sulphuric acid. The action is too slow to have practical value. 

Sulphur is oxidised by hot concentrated nitric acid, with bromine 
as the best catalyst, to sulphuric acid. 

S + 6HNO3 -»■ H a SO« + 6NO, + 2H t O 

This reaction is fully discussed on p. 438. 


The following experiments show that sulphur exists in several 
different forms, called 'allotropes'. The meaning of this term will be 
considered more fully after the experiments have been described. 

Preparation of rhombic or octahedral sulphur (a-sulphur) 

Shake some powdered sulphur with carbon disulphide for some 
time in a test-tube. (Take care to extinguish all flames in the vicinity.) 

Filter the contents of the test-tube into a dry beaker through a 
dry filter-paper and funnel. Fasten a filter-paper 
over the mouth of the beaker, pierce a few pin- 
holes in it, and set the beaker aside. The carbon 
disulphide will slowly evaporate, depositing 
crystals of sulphur, which, because of the slow 
evaporation, will be large enough for their shape 
to be seen. They will have the shape shown in 
Fig. 124. 

This variety of sulphur is called rhombic sul- 
phur or octahedral sulphur or a-sulphur. 

Note especially that the formation of the 
crystals takes place at ordinary room tempera- 

Fio. 124. 

Crystal of rhombic 



Preparation of monoclinic or prismatic sulphur (f3 -sulphur) 

Place powdered sulphur in a very large crucible or an evaporating 
dish. Heat it and stir, gradually adding more sulphur until the 
crucible or dish is almost brim-full of molten sulphur. Use a small 


Single crystal 

Fio. 125. 
Monoclinic sulphur. 

flame for the heating or the sulphur 
may begin to burn. Then allow 
the sulphur to cool. After a time, 
a solid crust will begin to form 
on the surface. When the crust is 
continuous, pierce it at two widely 
separated points with a glass rod 
and rapidly pour out the liquid 
sulphur from inside. With a pen- 
knife, cut through the solid crust 
all the way round the crucible or 
dish, near the rim, and lift it out. 
Underneath will be seen long 
'needle-shaped' crystals of sulphur 
whose shape is shown in Fig. 125. 
They are crystals of monoclinic 
sulphur or prismatic sulphur or (i-sulphur. 

Note that this variety crystallises in close contact with hot, molten 

Preparation of Amorphous Sulphur (8-sulphur) 

This variety of sulphur may be prepared in several ways. One is to 
saturate distilled water with hydrogen sulphide and then expose the 
solution to the air. Sulphur is deposited as an almost white powder, 
amorphous sulphur or b-sulphur. 

2H.S + O t ->-2H s O + 2S 

A llotropy and allotropes 

In the experiment just described, we have prepared three different 
varieties of the element sulphur. They have different properties, e.g., 
their densities (g/cm») differ (rhombic, 2.08; monoclinic, 1.98), but 
they all consist of pure sulphur and nothing else. When an element 
can exist in several different forms in this way it is said to show 
allotropy. (See p. 287.) 

Relation between monoclinic and rhombic sulphur 

The factor determining which of these two allotropes will be 
obtained in an experiment is temperature. In our experiments, rhom- 
bic sulphur was crystallised by evaporation of a solution of sulphur 
in carbon disulphide at ordinary room temperature, while mono- 
clinic sulphur was crystallised in contact with a mass of hot, molten 
sulphur. Roughly, then, we may say that if the sulphur crystallises 
while still hot, it does so as the monoclinic allotrope; if it crystallises 
while cold, the rhombic allotrope is formed. 



We can go further. Experiment has shown that the temperature 
which separates the two varieties is 96°C. If sulphur crystallises 
above this temperature, monoclinic crystals are formed, and if below 
it, rhombic. This temperature, 96°C is therefore called the 'transition 
temperature' between the two varieties. 

If rhombic sulphur, stable below 96°C, is kept above that tem- 
perature, it changes its crystalline form and becomes monoclinic, 
while, if monoclinic sulphur, stable above 96°C, is kept below that 
temperature, it slowly yields rhombic sulphur. 

Formation of plastic sulphur 

Heat some powdered roll sulphur in a test-tube until it is boiling 
rapidly. (The changes which occur are fully considered on p. 390.) 
Then pour the boiling sulphur in a thin continuous stream into a 
beaker full of cold water. It forms long, elastic, light-yellow ribbons 
of 'plastic sulphur', which are insoluble in carbon disulphide. This 
variety is not a true allotrope of sulphur. If kept for a few days, 
plastic sulphur becomes hard. This hard variety of sulphur is in- 
soluble in carbon disulphide. 

Comparison of two alio tropes of sulphur 

Rhombic (octahedral) Monoclinic (prismatic) 

Yellow translucent crystals. 
Density 2.08 g/cm 8 . 
Melting-point 114°C. 
Stable at temperatures below 

Transparent amber crystals. 
Density 1.98 g/cm a . 
Melting-point 119°C. 
Unstable at temperatures below 

96°C, reverting to rhombic 


Experimental evidence of the chemical identity of these allotropes is 
given by the fact that each is convertible into the other (by tempera- 
ture change) without change of weight and, if equal weights of the 
allotropes are converted into a given compound (e.g., sulphur 
dioxide), identical weights of product are given. 

Questions on sulphur will be found on page 421. 

Chapter 30 

Hydrogen Sulphide and the Sulphides 


(also called Sulphuretted Hydrogen) 


HYDROGEN sulphide was obtained in the experiment described 
on p. 7 by the action of dilute hydrochloric acid on ferrous sul- 
phide (iron(II) sulphide). This is the most convenient method of 
preparation, using the apparatus shown in Fig. 126. 

..cone, hydrochtoric 
z acid 

iron (n) sulphide^ 
and water 


Warm water 

Fig. 126. 
Preparation of hydrogen sulphide. 

If it is required to prepare the gas starting from sulphur, the best 
way is to prepare iron(H) sulphide first by the method of p. 7 and 
then to use it in the way about to be described. 

As the acid reaches the iron(II) sulphide, effervescence begins and 
the hydrogen sulphide is collected over water. It is rather soluble in 
cold water (about three volumes of the gas in one volume of water), 
but like all gases, it is less soluble in hot water. 

FeS(c) + 2HCl(aq)-»- FeCl.(aq) + H,S(g) 


Dilute sulphuric acid may also be used. 

FeS(c) + H,SO«(aq)-»- FeS0 4 (aq) + H s S(g) 

Characteristic Test. Soak a strip of filter-paper in lead 
acetate solution and drop it into a gas-jar of hydrogen 
sulphide. The paper turns dark brown or black. This 
colour change is caused by precipitation of black lead(Il) 

PKQHaOOs + H a S -> PbS| + 2C.H40, 

A purer specimen of hydrogen sulphide may be obtained by warm- 
ing antimony sulphide with concentrated hydrochloric acid. 

Sb 2 S s + 6HC1 -> 2SbCI 3 + 3H a S 
Tf required dry, the gas may be dried by passing it over calcium 
chloride and collected by downward delivery, as the gas is somewhat 
denser than the air. 

Kipp's apparatus 

Kipp's apparatus is a device for 
obtaining intermittent supplies of a 
frequently used gas such as hydrogen, 
carbon dioxide or hydrogen sulphide. 
(See Fig. 127). 

When the tap A is opened, the 
acid rises into the bulb B and attacks 
the ferrous sulphide, producing hydro- 
gen sulphide, which is delivered 
through A. When the gas is no 
longer required, A is turned off. The 
gas is still being generated which 
raises the pressure in B. The acid is 
therefore forced out of B and up into 
C. The generation of hydrogen sul- 
phide now stops because acid and 
ferrous sulphide are no longer in con- 
tact and the apparatus will remain 
inactive until tap A is again opened to 
obtain gas. 



ferrous Sulphlda 

Fio. 127. 
Kipp's apparatus. 


Appearance. The gas is colourless. 

Smell. Hydrogen sulphide has a repulsive, rather sweet smell 
similar to that of a rotten egg. It is, in fact, given off from putrefying 
eggs and also from decaying cabbages, both of which contain sulphur. 

Though very poisonous, the gas gives ample warning of its 
presence by its powerful smell. 

Solubility in water. Invert a gas-jar of hydrogen sulphide in cold 
water, remove the cover and shake gently. The rise of the water shows 
that the gas is fairly soluble. At ordinary temperatures, one volume 
of water can dissolve about three volumes of hydrogen sulphide. 

To the solution in the gas-jar, add blue litmus solution. It is turned 
claret colour. The solution is weakly acidic. It is known as hydro- 
sulphuric acid, but the name is very seldom used. It is one of the 
weakest acids known. Hydrogen sulphide is contained in the water 
of the sulphur springs round which have grown such spas as Harro- 
gate. These 'waters' are said to have curative properties. They 
certainly possess, as a consequence of their hydrogen sulphide 
content, all the unpleasant taste usually associated with medicines! 
Density. The gas has a density of 17 compared with that of hydro- 
gen, and is somewhat denser than air, which is 14.4 times denser than 

Combustion of hydrogen sulphide with a plentiful supply of air 

When several gas-jars of hydrogen sulphide have been collected 
remove the delivery-tube and fix the tube as shown below (Fig. 128). 
Apply a lighted taper. The hydro- 
gen sulphide burns with a blue 
flame similar to that of sulphur. 
Lower the tube into a wide gas-jar, 
closing the mouth with a square of 
cardboard and, when the flame is 
extinguished, remove the tube and 
add a weak, pink solution of acidi- 
fied potassium permanganate. On 
shaking, the solution becomes 
colourless and remains clear. This 
test proves the presence of sulphur 


Fio. 128. 

Combustion of hydrogen sulphide 
in a plentiful supply of air. 

2H.S + 30. -* 2H 4 + 2SO a 

Combustion of hydrogen sulphide with a limited supply of air 

Cut down the air supply to the flame obtained in the last section by 
putting into it a crucible lid. 

After a few seconds, a yellow deposit of sulphur will be seen on the 
lid. The reduced oxygen supply cannot oxidise the gas completely 
and free sulphur is deposited. 

2H 2 S(g) + 0,(g) -> 2S(c) + 2H a 0(l) 




Hydrogen sulphide as a reducing agent 

Hydrogen sulphide is a powerful reducing agent as the following 
experiments show. Like all reducing agents, it operates as a supplier 
of electrons. The usual product is a precipitate of sulphur, arising 
from the changes: 

fH 8 S ^ 2H + + S 2 ~ 
is 2 - -*-S + 2e- 

The electrons are accepted by the oxidising agent with which the 
hydrogen sulphide is reacting. (Very powerful oxidation, e.g., by 
concentrated nitric acid, may convert hydrogen sulphide to sulphuric 

Action of hydrogen sulphide with nitric acid. Dilute some con- 
centrated nitric acid with about one-third of its volume of water in a 
boiling-tube and pass hydrogen sulphide into it. Brown fumes of 
nitrogen dioxide are given off, a pale yellow deposit of sulphur 
appears, and the liquid becomes hot. The hydrogen sulphide has 
reduced the nitric acid to nitrogen dioxide and has itself been 
oxidised to sulphur. 

2HNO a + H 2 S ->• 2H 8 + 2NO a + S j. 
The solution also contains sulphuric acid, produced by the reaction: 
H 8 S + 8HNO3 -»■ H 8 S0 4 + 8NO a + 4H 8 
Action of hydrogen sulphide with iron(III) chloride solution. Per- 
form the experiment as above, using iron(III) chloride solution. 

A yellow deposit of sulphur appears and, on heating to coagulate 
the sulphur and filtering, a pale green solution of iron(II) chloride is 
obtained. The hydrogen sulphide has reduced the yellow iron(IIT) 
chloride to green iron(Il) chloride, being itself oxidised to hydrogen 
chloride, which dissolves in the water, and sulphur. 
2FeCl 3 + H 8 S -> 2FeCl 8 + 2HCI + S 
or 2Fe !+ + S a - —*■ 2Fe 2f + SJ 

Action of air on hydrogen sulphide. Pass a stream of hydrogen 
sulphide into distilled water in a beaker for about half an hour. Leave 
the solution exposed to air. After a few days a white deposit of 
amorphous sulphur will have appeared. The oxygen of the air has 
oxidised the hydrogen sulphide to sulphur and water. 

2H 8 S + 8 ->-2H 8 + 2S 
Hydrogen sulphide will reduce concentrated sulphuric acid, 
depositing sulphur. For this reason, the acid cannot be used to dry 

3H 8 S + H„S0 4 -> 4H t O -f- 4S 
Acidified potassium permanganate and dichromate solutions are 



reduced by the gas. The effect diners from that produced by sulphur 
dioxide because, while either gas decolorises the permanganate and 
turns the dichromate from yellow to green, hydrogen sulphide leaves 
also a precipitate of sulphur, while sulphur dioxide does not. 

2KMnO« + 5H,S + 3H 8 S0 4 -»- K 8 S0 4 + 2MnSO« + 8H 8 + 5S| 
K 8 Cr 2 7 + 3H,S + 4H 8 SO«-*K 8 SO t +Cr 8 (S0 4 ) 3 +7H 8 0+3S| 

For the action of hydrogen sulphide with sulphur dioxide, see 
p. 406, and for its action with the halogen elements, pp. 361, 385. 

The action of hydrogen sulphide on salts of metals 

Copper{II) sulphate. Heat a solution of copper sulphate in a boiling- 
tube and pass hydrogen sulphide into it. A dark brown precipitate 
appears, copper(II) sulphide. Filter the mixture. If sufficient hydrogen 
sulphide has been passed, the filtrate will be colourless because all the 
copper, which formerly coloured it, is now precipitated as copper(II) 
sulphide. The filtrate is dilute sulphuric acid. 

CuS0 4 + H 8 S -> CuS + H 8 SO« 


Lead{ll) nitrate. Experiment as above. 

Here, a black precipitate of Iead(II) sulphide is produced and the 
filtrate is dilute nitric acid. 

Pb(NO s ), + H 8 S -► PbS + 2HNO, 

Zinc sulphate. Experiment as above. 

A white precipitate of zinc sulphide is left and the filtrate contains 
dilute sulphuric acid. 

ZnS0 4 + H 8 S -► ZnS + H 8 SO« 
These reactions are all examples of double decomposition (see p. 160). 

Hydrogen sulphide as an acid 

Hydrogen sulphide acts as a weak dibasic acid. It forms with 
sodium hydroxide two salts, normal sodium sulphide, Na 8 S, 

2NaOH + H 8 S -> Na 8 S + 2H 8 
or, with excess of hydrogen sulphide, the acid salt, sodium hydrogen 
sulphide, NaHS. 

NaOH + H 8 S -* NaHS + H t O 

Potassium hydroxide reacts similarly. 

With aqueous ammonia, the gas gives yellow ammonium sulphide, 
mainly NH 4 HS. This compound is used in qualitative analysis. 

NH 4 OH + H 8 S -* NH 4 HS + H 8 Q 



Laboratory preparation of sodium sulphide 

The possible reactions of hydrogen sulphide with sodium hydroxide 
solution are: 

2NaOH + H,S —*■ Na a S + 2H.O 
2NaOH + 2H.S — »• 2NaHS + 2H.O 

It is clear, from the equations, that the volume of hydrogen sulphide 
needed to convert a given weight of sodium hydroxide into sodium hydro- 
gen sulphide, NaHS, is twice that required to convert it to sodium sulphide, 
Na 3 S. It is impossible in practice to determine when just enough hydrogen 
sulphide has been used to convert the alkali into sodium sulphide, so the 
best way of carrying out the preparation is to convert half of the sodium 
hydroxide into sodium hydrogen sulphide by saturation with hydrogen 
sulphide, and then to form the normal salt from the acid salt by addition 
of the other half of the sodium hydroxide. 

NaOH + H,S -> NaHS + H.O 
NaHS + NaOH — >■ Na,S + H.O 
Experiment. Measure out 50cm'of bench(about 2 M) sodium hydroxide 
solution, divide it into two equal parts and, into one of them, pass hydrogen 
sulphide until no more is absorbed and the liquid smells strongly of 
hydrogen sulphide. Add the other half of the sodium hydroxide solution 
and obtain crystals of sodium sulphide by the method described on p. 230. 



Piece of Lead 
Acetate paper 
is blackened 

- \ Boiling 


Fio. 129. 
Proof of composition of hydrogen sulphide. 

Proof that hydrogen sulphide contains only hydrogen and sulphur 

Heat sulphur in a boiling-tube till it is boiling, then pass through it a 
stream of dry hydrogen by apparatus shown in Fig. 129. The wet lead 
acetate paper will be turned black. This proves that hydrogen sulphide is 
present in the gases. Since it can only have been produced by direct cora- 
bimuion of hydrogen and sulphur, it must contain these two elements 

H, +S— *-H,S 

This is not a practical preparation of hydrogen sulphide. Only a very small 
proportion of the hydrogen is converted to hydrogen sulphide. 


Formula of hydrogen sulphide 
This is considered on p. 1 19. 











Y Sulphides of these metals 
fare soluble in water. 

Sulphides of these metals 
are insoluble in water. 

Sulphides of these metals will 
[not precipitate from acidified 

Sulphides of these metals will 
\precipitate from acidified solu- 

Potassium sulphide, K 2 S 
This is similar to sodium sulphide and is similarly prepared. 

Sodium sulphide, Na,S 

The preparation of this compound by neutralisation of sodium 
hydroxide by hydrogen sulphide is described on p. 400. Its aqueous 
solution is alkaline and smells of hydrogen sulphide. When heated 
with sulphur it forms 'poly sulphides' of sodium. For example: 
Na 2 S + 4S-*Na 2 S B 

In industry, sodium sulphide is prepared by heating sodium 
sulphate with coke (p. 418). 

Uses. (1) For preparing a class of very 'fast' dyes. 
(2) For stripping the hair from hides. 

Calcium sulphide, CaS 

This compound was chiefly important in the form of the 'alkali 
waste' of the Leblanc process. Sulphur was recovered from it. 

If it contains traces of certain metals, for example, 0.01 % bismuth, 
it is 'phosphorescent', that is, after exposure to light, it will emit a 
violet glow whose intensity gradually diminishes. The glow fades out 
after some hours. 

Zinc sulphide, ZnS 

This compound occurs as the mineral 'zinc blende'. It may be pre- 
cipitated by hydrogen sulphide from a neutral (or alkaline) solution 
of a zinc salt. . 

Zn s+ + S»~ -»■ ZnS J 



Like calcium sulphide, and under similar conditions, zinc sulphide 
is phosphorescent. The luminous paint on watches is usually zinc 
sulphide, containing about 1 part of a radium salt in 100 000 000 of 
the sulphide. 

Iron(II) sulphide, FeS 

This black, insoluble compound is usually employed for the pre- 
paration of hydrogen sulphide (p. 395). It is prepared by heating iron 
with sulphur in the calculated quantities (p. 391). 

Iron disulphide (iron pyrites), FeS 2 

This occurs as a hard, brassy mineral. There are great masses of it 
in Spain. It is the cheapest source of sulphur dioxide, which it gives 
off when burnt in air (p. 41 1). 

Lead(II) sulphide, PbS 

Lead (II) sulphide occurs as the mineral galena, and is precipitated 
from solutions of lead salts by hydrogen sulphide (p. 399). The most 
satisfactory test for hydrogen sulphide (p. 396) is the production of a 
dark brown (almost black) stain of lead(II) sulphide on a filter-paper 
soaked in lead(II) acetate solution. 

Pb a+ + S 2 - ->- PbS | 

Copper(II) sulphide, CuS 

This is a black insoluble compound precipitated from a solution of 
a copperfH) salt by hydrogen sulphide (p. 399). 

Cu»+ + S a " -* CuS | 
Questions on sulphides will be found on page 421. 

Chapter 31 

Oxides and Acids of Sulphur 
and the Sulphates 



THIS compound, which is a gas under ordinary conditions, is con- 
veniently prepared in the laboratory by the apparatus of Fig. 130. 
There is no action until the mixture in the flask becomes hot Then 
rapid effervescence occurs and the sulphur dioxide, being very soluble 
in water and denser than air, is usually collected as shown. It may 
also be collected over mercury. 

Cu(c) + 2H 8 S0 4 (1) -»• CuS0 4 (c) + 2H,0 + SO t (g) 

, Sulphuric 
' Acid 


Cone. Sulphuric acid 
to dry the gas 

Fio. 130. 

Preparation of sulphur dioxide. 




A dark brown mixture is left in the flask. It contains anhydrous 
copper(II) sulphate and certain impurities. Crystals of copper(II) 
sulphate may be obtained from it by the method described on p. 231. 
Tests. 1. The gas has a very irritating smell and a metallic 

2. Action on potassium permanganate solution. The solu- 
tion is turned from purple to colourless by sulphur dioxide. 
(No precipitate is left as in the case of reduction of the per- 
manganate by hydrogen sulphide.) 

5SO, + 2KMnO t + 2H a O -> K,S0 4 + 2MnS0 4 + 2H,S0 4 
The explanation of the change in colour is that the potassium 
permanganate is decomposed and all the products of the reaction 
give colourless solutions. (The manganese sulphate is too small in 
amount for its very pale pink colour to be observed.) 

Properties of sulphur dioxide 

Appearance. The gas is colourless. 

Smell. The gas has an irritating smell and a rather sweet taste. It 
is fairly poisonous and is used for fumigation. 

Solubility in water. Invert a gas-jar of the gas in cold water and 
shake. The rapid rise of water shows that the gas is readily soluble in 
water. Add to the liquid some blue litmus solution. It is turned red. 
The solution is acidic. 

The sulphur dioxide reacts chemically with the water to produce 
sulphurous acid. 

H.O + SO a ?± H a SO g *± 2H+ + S0 8 «- 
This acid will be considered more fully later. 

Sulphur dioxide as a reducing agent 

Sulphur dioxide, in the presence of water, is a powerful reducing 
agent. It reacts with water to form sulphurous acid and the sulphite 
ion, S0 3 * _ , and this ion, like reducing agents in general, acts as a 
supplier of electrons. This occurs in association with water. 
H s O + SO, ^ H 2 SO, ^ 2H + + S0 3 2 ~ 
SCy- + H.O -► S0 4 »- + 2H+ + 2e~ 
The electrons are accepted by the oxidising agent with which the 
S0 2 -water system is reacting, e.g., 

iron(III) ion, which is reduced to iron(II) ion : Fe 3+ + e~—*- Fe 2+ 
chlorine, which is reduced to its ions: Cl a + 2e~ —*■ 2C1 - 
acidified potassium permanganate, which is reduced to a man- 
ganese(II) salt: 

Mn0 4 " + 8H+ + 5e~ -*■ Mn* + + 4H g O 



The following are important examples of the reducing action of sul- 
phur dioxide in aqueous solution. 

(a) Action of sulphur dioxide with concentrated nitric acid. Put some 
concentrated nitric acid into a boiling-tube and pass into it a current of 
sulphur dioxide from a siphon of liquid sulphur dioxide. Brown fumes 
are evolved (nitrogen dioxide) and the liquid becomes warm. Dilute 
some of the liquid and add dilute hydrochloric acid and barium chlor- 
ide solution (the recognised test for a soluble sulphate). The white 
precipitate of barium sulphate proves the presence of sulphuric acid. 

BaCl, + H 8 S0 4 -»- BaS0 4 + 2HC1 

The concentrated nitric acid has oxidised the sulphur dioxide in 
the presence of water to sulphuric acid and has been itself reduced to 
nitrogen dioxide. 

SO» + 2HN0 3 -»- H,,S0 4 + 2N0 2 

(b) Action of sulphur dioxide on iron(HI) sulphate solution. Make a 
solution of iron(III) sulphate (or iron ammonium alum) in water in a 
boiling-tube and pass into it sulphur dioxide as above. The brownish 
colour of the solution is rapidly converted to pale green. 1 The sulphur 
dioxide has reduced the brown iron(III) sulphate to light green 
iron(II) sulphate and has itself been oxidised to sulphuric acid. 

2IV+ + 2H s O + S0 2 -► 2Fe*+ + S0 4 °- + 4H + 

(c) Action of sulphur dioxide on potassium dichromate. Acidify a 
solution of potassium dichromate in a boiling-tube with dilute sul- 
phuric acid, and pass through it a stream of sulphur dioxide from a 
siphon of liquid sulphur dioxide. There is a rapid colour change from 
golden yellow to green, but no precipitate appears (compare the 
action of hydrogen sulphide, p. 399). 

K e Cr 2 7 + 3S0 3 + H 2 S0 4 -► K 3 S0 4 + Cr^SOJ, + H t O 

The potassium dichromate has oxidised the sulphur dioxide in the 
presence of water to sulphuric acid, being itself reduced to green 
chromium(III) sulphate. 

(d) Action of sulphur dioxide on potassium permanganate. (See test 
for sulphur dioxide p. 404.) 

The potassium permanganate oxidised the sulphur dioxide in the 
presence of water to sulphuric acid, and was itself reduced to man- 
ganese(II) sulphate. 

5SO t + 2KMnO« + 2H s O ->• K«S0 4 + 2MnS0 4 + 2H,S0 4 

1 The red solution which may be formed is a complex sulphite which de- 
composes on heating, leaving the products as indicated by the equation above. 



Bleaching action of sulphur dioxide 

Sulphurous acid is a bleaching agent. This may easily be shown by 
dropping into a gas-jar of the gas (containing some water) a few blue 
flowers, e.g., blue crocus, iris, or bluebells. After a few minutes, the 
flowers will have lost their blue colour. 

This bleaching is also a reducing action. The sulphurous acid takes 
up oxygen from the colouring matter of the flowers and forms sul- 
phuric acid; the removal of oxygen from the dye converts it to a 
colourless compound. Sulphur dioxide is used industrially for bleach- 
ing sponges and straw for straw hats. The oxygen of the air may 
oxidise the reduced colourless compound back to the original coloured 
compound, which explains why straw hats gradually become yellow 
with use. 

Action of sulphur dioxide on halogen elements. See p. 363. 

Action of sulphur dioxide with hydrogen sulphide 

Add to a gas-jar of sulphur dioxide a little water, invert over it a 
gas-jar of hydrogen sulphide and allow the gases to mix. A yellow 
deposit of sulphur will be produced at once. The dry gases do not 

2H 2 S + S0 8 -*-2H s O + 3S 
Note that, here, the sulphur dioxide is actually acting as an oxidising 
agent, supplying oxygen to the hydrogen sulphide. As we have seen 
above, however, sulphur dioxide usually shows reducing properties. 
Here it has encountered in hydrogen sulphide a more powerful 
reducer than itself, which takes up its oxygen and causes it to act as 
an oxidiser. 

Gass experiments confirming properties of sulphur dioxide 

Put one or two grams of sodium sulphite crystals into a test-tube, 
cover them with dilute hydrochloric acid and warm. Dip a clean glass 
tube into a very dilute solution of potassium permanganate and lower 
it into the gas. Be careful not to lower it into the liquid or into the 
spray immediately above the liquid. The pink colour is discharged. 
Remove the tube, wash well with water and repeat using the following: 

(a) potassium dichromate solution, 

(b) blue litmus solution, 

(c) chlorine water, 

(d) hydrogen sulphide water, 

(e) barium chloride solution, 

(/) barium chloride solution acidified with dilute hydrochloric 
(Sec pages 404, 405, for explanations.) 



Action oflead(lV) oxide on sulphur dioxide. Warm some lead(lV) 
oxide (lead dioxide) on a deflagrating spoon and lower it into a gas- 
jar of sulphur dioxide. The lead dioxide glows and a white deposit of 
lead(II) sulphate is left. 

Pb0 2 + SO t -*-PbS0 4 

Liquefaction of sulphur dioxide 

Sulphur dioxide can readily be liquefied by being dried by con- 
centrated sulphuric acid and passed through a freezing mixture of ice 
and salt. It liquefies under ordinary atmospheric pressure at about 
— 10°C. It can be kept liquid at ordinary room temperature if under 
slight pressure, and it is sold in siphons under pressure. 

Formula of sulphur dioxide. See p. 120. 

Sulphur dioxide in chemical industry 

Sulphur dioxide is very important as an intermediate compound 
in the manufacture of sulphuric acid (p. 411). It is prepared by burn- 
ing sulphur in air. 

s + o,->so. 

or by burning iron pyrites in air. 

4FeS i! + 110, -»• 2Fe t 3 + 8SO t 


This acid has never been obtained free from water. Any attempt to 
prepare the pure acid always results in its decomposition into sulphur 
dioxide and water. 

It is prepared by passing sulphur dioxide into water. The gas is 
readily soluble and it is advisable to prevent 'sucking back' by the 
use of a funnel just touching the water surface (Fig. 131). 

The reaction in the flask is the same as described under the pre- 
paration of sulphur dioxide (p. 403). It is, of course, not necessary 
here to dry the sulphur dioxide. 

Cu + 2H,S0 4 -»- CuS0 4 + 2H t O + SO, 

Sulphur dioxide is the anhydride of sulphurous acid and may be 
called 'sulphurous anhydride'. 

Definition. An anhydride is the oxide of a non-metal, which, when com- 
bined with water, forms an acid. 

CO, + H,0 s» H,CO, ^ 2H+ + C0 3 8 - 
carbonic carbonic 

anhydride acid 



SO, -|- H 2 ^ H 2 SO s ^ 2H+ + S0 3 »- 
sulphurous sulphurous 

anhydride acid 

SO„ + H 2 ^ H 2 S0 4 ^ 2H + + S0 4 *- 

sulphuric sulphuric 

anhydride acid 

An anhydride will not always combine directly with water to give the 
corresponding acid, e.g., silicon dioxide, Si0 2 , is the anhydride of 
silicic acid, H 2 Si0 3 , though the acid cannot be prepared by direct 

| j Sulphuric 

Fio. 131. 
Preparation of sulphurous acid. 

combination of its anhydride with water. The acid is prepared from 
one of its salts and, when heated, loses water, leaving silicon dioxide 
as the residue. 

Properties of sulphurous acid 

Sulphurous acid is a colourless liquid which smells strongly of 
sulphur dioxide. 

Reducing action 

(0 The acid has all the reducing actions described previously as 
those of sulphur dioxide in the presence of water (pp. 404-5). 

(ii) Effect of exposure to air. Leave a beaker of sulphurous acid 
exposed to air for a few days. Then add to it hydrochloric acid and 
barium chloride solution. The white precipitate of barium sulphate 



proves that the oxygen of the air has oxidised the sulphurous acid to 

sulphuric acid. 

2H 3 S0 3 + 2 ->2H 2 S0 4 
Ba"(aq) + S0 4 2 "(aq) -> BaS0 4 (c) 

Action of sulphurous acid with alkalis 

Sulphurous acid is a dibasic acid and with sodium hydroxide forms 
two sodium salts, the acid salt, sodium hydrogen sulphite, NaHS0 3 , 
and the normal salt, sodium sulphite, Na a S0 3 . 

NaOH + H 2 S0 3 -> NaHSO s + H s O 
2NaOH + H 3 S0 3 -»- Na 2 S0 3 + 2H 2 
Potassium hydroxide solution behaves similarly. 

Laboratory preparation of sodium sulphite. This is similar to the 
preparation of sodium sulphide, described on p. 400, using sulphur 
dioxide instead of hydrogen sulphide. 

NaOH + H 2 SO a ->- NaHS0 3 + H a O 
NaHS0 3 + NaOH -* Na 2 S0 3 + H 2 
Sulphites give off sulphur dioxide when warmed with dilute hydro- 
chloric acid or dilute sulphuric acid (test for S0 2 , p. 404), e.g., 

Na 2 SO a + H 2 S0 4 -*■ Na 2 SO« + H 2 + S0 2 
This is occasionally used as a method of preparing sulphur dioxide. 
Used in dilute acidified solutions, sulphites have all the reducing 
actions of sulphur dioxide and water, or sulphurous acid (see p. 404). 

Preparation of sulphurous acid, H 2 S0 3 , from sulphur 

To convert sulphur into sulphurous acid it is necessary first to 
oxidise the sulphur to sulphur dioxide and then absorb this gas in 

This can be done by the apparatus of Fig. 132. 



To filter 

W Sulphur Jj 


Fio. 132. 
Preparation of sulphurous acid from sulphur. 



Heat the sulphur and, by means of a filter-pump, draw over it a 
rapid stream of air. The sulphur burns and the sulphur dioxide pro- 
duced is absorbed as it passes through the water in the WoulfTs 
bottle. The liquid left is sulphurous acid. Sulphur vapour may be 
earned over unburnt and appear as a yellow precipitate in the bottle. 
Remove it by filtration. 

S + 8 ->SO, 
H t O + S0 3 ->H 3 S0 8 

The method of producing sulphurous acid given on p. 403 is much 
more convenient in the laboratory, but the sulphurous acid prepared 
on the large scale is made by modification of the above method or 
by burning iron pyrites, FeS a . 

4FeS 3 + HO, -»- 2Fe 3 O s + 8SO, 


This compound is a white hygroscopic solid. A sample of it is 
usually kept in a sealed glass bulb as a laboratory exhibit. 

Preparation of sulphur trioxide 

It is prepared by passing a mixture of dry sulphur dioxide and dry 
air, or oxygen, over heated platinised asbestos (or vanadium pent- 
oxide). Platinised asbestos is made by soaking asbestos in platinum 
chloride solution and then igniting it, when platinum is left in a very 
finely divided form. 

PtCI 4 ->- 2C1, + Pt 



platinised asbestos 

— -feT 

acid to 

dry the 


Wide tube 


mixture | 



Fig. 133. 
Preparation of sulphur trioxide. 



The platinum is a catalyst and the best temperature is 45O°-500°C. 
The sulphur trioxide is seen as dense white fumes and may be solidi- 
fied in a freezing mixture of ice and a little common salt (Fig. 133). 

2S0 3 +0»-»-2S0 3 
The sulphur trioxide container is protected from atmosphericmoisture 
by a calcium chloride tube. Sulphur trioxide is important because it 
combines vigorously with water, giving sulphuric acid. 

H 3 + S0 3 -»-H a S0 4 
It is the anhydride of this acid and sulphur trioxide may be termed 
'sulphuric anhydride'. (Sec also p. 408.) 


Lead chamber process for the manufacture of sulphuric acid 

To offset Britain's lack of the raw materials, sulphur and pyrites, 
a new process is producing sulphur dioxide by strongly heating 
anhydrite, CaS0 4 , with coke. The main reaction is: 

2CaS0 4 + C -> 2CaO + 2SO, + CO a 
By including sand and ashes containing alumina, the quicklime is 
converted to a valuable by-product, cement (calcium silicate and 
aluminate). Pyrites is also burnt in air to produce sulphur dioxide. 

4FeS a + 110 8 -*• 2Fe a 3 + 8S0 3 
It is converted by the oxygen of the air, in the presence of steam, into 
sulphuric acid. Nitrogen monoxide is used as a catalyst or oxygen- 

2NO-f 3 ->-2N0 3 

N0 2 + H 3 + SO, -* NO + H 2 S0 4 

The nitrogen monoxide is usually supplied in a modern plant by 
oxidising ammonia by oxygen of the air. The two gases are passed 
over heated platinum. 

4NH 3 + 50 2 -► 4NO + 6H 2 
The main oxidation of the sulphur dioxide is carried out in large lead 
chambers, on the floors of which 'chamber-acid' (65% sulphuric 
acid) accumulates. It is not very pure but finds a ready sale where an 
acid of high purity is not needed. The plant used is made more 
elaborate by devices for recovery of the nitrogen monoxide, which 
would otherwise escape, and for its restoration into the reacting 



Contact process for manufacture of sulphuric acid (see also p. 157) 

Sulphur dioxide (prepared by burning sulphur) and air are passed 
oyer a catalyst, heated to 450°-500°C. About 98% of the possible 
yield of sulphur trioxide is obtained. 

2SO t + O t ^ 2SO, 
Originally, platinised asbestos was used as catalyst but platinum is 
very expensive and easily 'poisoned' by impurity which made elabor- 
ate purification of the gases necessary (especially from arsenical com- 
pounds). Vanadium pentoxide, V 2 5 , has replaced platinum as the 
usual catalyst employed. 

The sulphur trioxide cannot be satisfactorily absorbed by water. 
A mist of fine drops of dilute sulphuric acid fills the factory if direct 
absorption in water is tried. It is dissolved in concentrated sulphuric 
acid, forming a fuming liquid called 'oleum' for which there is some 
demand. Most of the 'oleum' is carefully diluted with the correct 
amount of water to give ordinary concentrated sulphuric acid. 
SO a + H 2 -»- H 2 S0 4 

Properties of sulphuric acid 

Sulphuric acid is a dense oily liquid, 'Oil of Vitriol'. It has several 
very important properties. 

Dilute sulphuric acid — as an acid 

Sulphuric acid is dibasic, ionising in two stages to produce first a 
hydrogen ion and a bisulphate or hydrogen sulphate ion, HS0 4 - , from 
a molecule of the acid, after which the bisulphate ion may ionise 
further to produce a hydrogen ion and the sulphate ion, S0 4 2 ~ 
H 2 S0 4 ^ H+ + HS0 4 - ^ 2H + + S0 4 2 ~ 

In dilute solution, the acid is almost completely ionised and so is a 
strong acid. 

Because of its dibasic character, this acid forms two sodium salts, 
sodium sulphate, (Na') 2 S0 4 2 -, and sodium hydrogen sulphate or 
sodium bisulphate, Na + HS0 4 - . 

2(Na+OH-) + (H+) a S0 4 *~ -»- (Na+) 2 S0 4 2 - + 2H g O 
Na'OH- + (H+) 2 S0 4 s - ->- Na + HS0 4 " + H 2 

Preparation of sodium sulphate and hydrogen sulphate. It is evident from 
the previous equations that the amount of sodium hydroxide needed to 
convert a given amount of sulphuric acid into sodium hydrogen sulphate is 
half that required to convert it to sodium sulphate. 

Measure out, say, 100 cm" of bench (2 M) sodium hydroxide solution 
into a flask, add litmus and then run in carefully, from a burette, bench 
(2 M) dilute sulphuric acid, until the solution is neutral (purple). Note the 
volume of dilute sulphuric acid needed (say x cm*). This solution now 
contains sodium sulphate. 



2NaOH + H,S0 4 — »-Na,S0 4 + 2H.O 
Then measure out a further 1 00 cm 3 of the same sodium hydroxide solution 
and add to it, from the burette, 2x cm" of the same acid. This solution now 
contains sodium hydrogen sulphate. 

2NaOH + 2H.SO, -> 2NaHS0 4 + 2H,0 
Obtain crystals in the usual way from both solutions (see p. 231). 

Similarly, two potassium salts, potassium sulphate, K,S0 4 , and potas- 
sium hydrogen sulphate, KHS0 4 , can be made. 

Dilute sulphuric acid also neutralises basic oxides or hydroxides 
to form salts and water, e.g., 

CuO + H 2 S0 4 -► CuS0 4 + H,0 
ZnO + H 2 S0 4 -► ZnS0 4 + H 2 
Cu(OH) 2 + H 2 S0 4 -> CuS0 4 + 2H 2 
Zn(OH) a + H 2 S0 4 -»• ZnS0 4 + 2H 2 

Action of dilute sulphuric acid with metals. Some of the common 
metals displace hydrogen from dilute sulphuric acid, e.g., 

Zn + 2H+ -»- Zn 2+ + H g 

Fe + 2H + -»-Fe 4+ -r-H 2 

Mg + 2H+ -»• Mg a+ + H» 

Copper is, however, without action on this acid. Note that cold, 
concentrated sulphuric acid, in the complete absence of water, is not 
attacked by any metal. 

Action of sulphuric acid with carbonates. If the sulphate of a metal 
is soluble, dilute sulphuric acid readily attacks its carbonate with 
evolution of carbon dioxide, e.g., 

Na t CO s + H 2 S0 4 -> Na 2 S0 4 + H g O + CO a 
MgCOa + H,S0 4 -► MgS0 4 + H 2 + CO, 
or CO s a - + 2H + -> H a O + C0 2 

If dilute sulphuric acid is added to marble, CaCO s , however, the 
effervescence is checked after a few seconds. This is because the 
calcium sulphate which is formed is only sparingly soluble in water 
and soon forms a deposit on the surface of the marble, separating it 
from the acid and checking the action. 

Concentrated sulphuric acid as an oxidising agent 

Like all oxidising agents, sulphuric acid acts as an acceptor of 
electrons. When hot and concentrated, the acid shows, as its principal 

2H 2 S0 4 + 2e~ -»• S0 4 *- + 2H s O + SO, 
The electrons are supplied by the reducing agent concerned in the 
reaction. This may be a metal, such as copper or zinc: 
Cu (or Zn) -> Cu 2+ (or Zn 2 -) + 2e~ 



The metallic ion is left associated with the S0 4 2_ ion as the correspond- 
ing metallic sulphate, and the reaction is usually written in a single 
equation, as: 

Cu (or Zn) + 2H 2 S0 4 -* CuS0 4 (or ZnS0 4 ) -j- 2H g O + SO, 

The non-metals, carbon and sulphur, are also oxidised by the hot, 
concentrated acid to give sulphur dioxide or carbon dioxide. 

S + 2H,S0 4 -* 2H 2 + 3S0 2 

C + 2H 2 S0 4 -*■ 2H 2 + 2S0 2 + C0 2 

The sulphur dioxide given off may be detected by the decolorisa- 
tion of potassium permanganate solution. 

Concentrated sulphuric acid possesses an affinity for water 

The acid has a very great affinity for water. It mixes with water 
with a very great evolution of heat. The two, when mixed in equal 
volumes at room temperature, may give a liquid whose temperature 
is as high as 120°C. This indicates chemical reaction but its nature is 
not clearly understood. It is very important when mixing the acid with 
water to add the acid to the water and NEVER the water to the acid. 
It is necessary to stir the liquid as the acid enters to prevent formation 
of a lower layer of acid. 

Concentrated sulphuric acid is hygroscopic, i.e., it absorbs water- 
vapour out of the air, increasing in bulk and becoming dilute. This 


acid \^ 



Exposure lo air 
just beginning. 

After exposure to air 
for about three weeks. 

Fio. 134. 
Exposure of concentrated sulphuric acid to the air. 

can be shown by the following experiment for which Fig. 134 is 
sufficient explanation. The concentrated acid is used for drying gases, 
e.g., sulphur dioxide, chlorine, hydrogen chloride. It cannot be used 
to dry a reducing gas like hydrogen sulphide, or an alkaline gas like 

So great is the affinity of concentrated sulphuric acid for water that 
it can decompose many compounds by removing from them the 
hydrogen and oxygen necessary to form water, with which it then 
combines. This is called a dehydrating action. 



Concentrated sulphuric acid as a dehydrating agent 

1. With sugar. 

Place about a tablespoonful of sugar in a 450 cm 3 beaker and cover 
it with water. Place the beaker in a trough, for safety, and pour in a 
steady stream of concentrated sulphuric acid. The sugar is charred 
and a spongy black mass of charcoal rises, filling the beaker. Steam 
is given off and the whole mass becomes very hot. 

The acid has taken out the elements of water from the sugar leaving 
a black mass of carbon. 

C„H m O„ (+ «H 2 SO«)->. 12C + (11H 2 + nUJOJ 

A similar action is the explanation of the very marked corrosive 
action of the acid on cloth, e.g., cotton. This is cellulose, whose 
simplest formula is C 4 H 10 O 6 . As above, 

C,H 10 O 5 (+ «H 2 S0 4 ) -► 6C + (5H 2 + nH 2 SO«) 
and a hole appears in the cloth. Similar reactions account for its 
rapid and serious burning of the skin. 

2. With oxalic acid, H 2 C 2 4 . 

Place a little oxalic acid in a test-tube, add a little concentrated 
sulphuric acid and warm gently. Effervescence occurs. Apply a lighted 
splint to the test-tube. The gas burns with a blue flame, showing that 
carbon monoxide is given off. Extinguish the flame and pass the gas 
into lime-water held in a boiling-tube. The turbidity shows that 
carbon dioxide is also present. The reaction is of the same type as 
those above, and it is used for the laboratory preparation of carbon 
monoxide (see p. 299). 

H 2 C 2 4 (+ H 2 S04) -> CO + CO, + (H a O + H.SOJ 

Uses of sulphuric acid 

In quantitative order, the uses of sulphuric acid in the United 
Kingdom (1968) were as given below. 
34% for fertilisers, such as ammonium sulphate and calcium 

15% in the manufacture of paints and pigments; 
13% in connection with natural and man-made fibres; 
12% for the production of other chemicals, such as metallic sul- 
phates, hydrochloric acid, hydrofluoric acid and plastics; 
9% in the manufacture of detergents and soap; 
4% in connection with the extraction and use of metals, in- 
cluding 'pickling' to clean metallic surfaces. 

The remaining 13% covered various minor uses. 


Test for sulphuric acid and soluble sulphates. To a little 
dilute sulphuric acid in a boiling-tube add dilute hydro- 
chloric acid and barium chloride solution. A white precipi- 
tate of barium sulphate is formed. This is the characteristic 
test for any soluble sulphate. 

BaCl a + H 2 S0 4 -»- 2HC1 + BaSO« J 


Methods of preparation 

These are fully dealt with in the chapter, 'Acids, Bases and Salts', 
pp. 222-37. Briefly summarised, they arc : 

(1) By the action of sulphuric acid on a metal. 

Dilute acid 

Fe + H 2 S0 4 ->• FeS0 4 + H, 

Mg + H 2 S0 4 -»- MgS0 4 + H, 

Zn + H s S0 4 -> ZnSO« + H, 

Hot concentrated acid 

Cu + 2H 2 S0 4 -*■ CuS0 4 + 2H 2 + SO, 

(2) By the action of dilute sulphuric acid on the oxide, hydroxide or 
carbonate of the metal. 

For example: CuO + H 2 S0 4 
Zn(OH) 2 + H 2 S0 4 
Na 2 C0 3 + H 2 S0 4 



CuS0 4 + H 2 
ZnS0 4 + 2H s O 

Na 2 S0 4 + H a O + C0 2 


(3) By double decomposition 

This is limited in application to the preparation of insoluble sul- 
phates. Only two common sulphates are insoluble, barium sulphate 
and lead sulphate. (Calcium sulphate is sparingly soluble.) 

Pb ,+ (aq) + S0 4 z -(aq)-* PbS0 4 (c) 
Ba*+(aq) + S0 4 *-(aq) -*■ BaS0 4 (c) 

Aluminium sulphate, Al.,(S0 1 ) 3 .18H..O 

This is a white solid. It is conveniently prepared by dissolving the 
oxide or hydroxide of the metal in dilute sulphuric acid, leaving the 

acid in slight excess to counter hydrolysis. The sulphate can be ob- 
tained by evaporation to small bulk and cooling, but it does not 
crystallise well. 

2A1(0H) 3 + 3H 2 S0 4 -> AIjCSOJs + 6H a O 
It is most commonly encountered in the form of potash alum, one 
of an important group of salts called the alums. 

The alums 
These are double salts of general formula 

X,S0 4 .Y a (S0 4 ) 3 .24H 2 or X + Y» + (S(V-) 2 .12HaO 

where X is Na, K or NH 4 
and Y is Fe(IU), Al or Cr. 1 

(Note that X is a monovalent and Y a /r/valent metal.) 
The alums crystallise well from water. They all have similar crystal- 
line shape, consequently crystalline layers of different alums may be 
deposited on one another to produce large, composite crystals with 
layers of varying colours. 

The two commonest alums are: 

Potash alum K 2 S0 4 .A1 2 (S04),.24H,0 (colourless) 
Iron(III) alum (NH 4 ) 2 S0 4 .Fe t (S0 4 ) 3 .24H 2 (purple) 

Preparation of potash alum 

(This is commonly called simply 'alum'). Potassium sulphate and 
aluminium sulphate are weighed out approximately in the propor- 
tions of their molecular weights 

K 2 S0 4 : A1 1 (S04),.18H 1 

174 g 666 g 

Use, say, one-twentieth of these figures, i.e., 8.7 g and 33 g. These 
amounts are dissolved, with heat, in as little water as possible. In the 
case of the aluminium salt, the water should be slightly acidified with 
dilute sulphuric acid. The hot solutions are then mixed and stirred. 
On cooling, colourless alum crystals separate out and are filtered, 
washed with cold distilled water and dried. 

Ammonium sulphate (sulphate of ammonia), (NH 4 ) 2 S0 4 

This compound is very widely used as a nitrogenous fertiliser. It 
may be made in the laboratory by neutralisation of dilute sulphuric 
acid with ammonia (p. 234). In industry, it is produced by the action 

1 Alums are also formed from some less common metals, for example caesium 
and rubidium. 



of ammonia and carbon dioxide on the mineral 'anhydrite', calcium 
sulphate, in the presence of hot water. 

CaS0 4 + 2NH 3 + CO a + H a O ->- CaC0 3 j + (NH 4 ) 2 S0 4 
The chalk is filtered off and the ammonium sulphate crystallised. 

Potassium sulphate, K.S0 4 

This compound may be prepared in the laboratory by neutralisa- 
tion of potassium hydroxide solution by dilute sulphuric acid (p. 234). 
In industry, it is usually prepared by heating potassium chloride with 
concentrated sulphuric acid. 

2KCI(c) + H 8 SO 4 0) -»■ K 2 S0 4 (c) + 2HCI(g) 

Unlike most soluble sulphates, it crystallises without water of crystal- 

Sodium sulphate, Na.SO., 

This salt is usually met with in the form of transparent crystals of 
the decahydrate, Na 2 SO 4 .10H 2 O, Glauber's salt. In the laboratory, 
it may be made by neutralising sodium hydroxide solution by dilute 
sulphuric acid (p. 234). In industry, it is prepared by heating sodium 
chloride with concentrated sulphuric acid. 

It is used in medicine, in the manufacture of glass, and, by heating 
with coke, for the manufacture of sodium sulphide. 

Na,S0 4 + 4C -> Na a S + 4CO 

Calcium sulphate, CaS0 4 

This salt occurs naturally as anhydrite, CaS0 4 , and gypsum, 
CaS0 4 .2H a O. For the use of anhydrite in making sulphuric acid and 
ammonium sulphate, see pp. 411 and 418. 

Gypsum is chiefly employed for the manufacture of plaster of 

Plaster of Paris (CaS0 4 )„.H 2 0, calcium sulphate hemihydrate. This 
compound is made by heating gypsum in large steel vessels of several 
tons capacity. The gypsum is stirred mechanically and the tempera- 
ture is maintained between 100°C and 200°C. 

2(CaS0 4 .2H 2 0)-v (CaSOJj.HjO + 3H 2 
plaster of Paris 

When mixed with water, plaster of Paris sets to a hard interlacing 
mass of fine needles of gypsum, expanding at the same time. It is 
used for making casts for statuary (the expansion during setting 
ensures a fine impression), in surgery to maintain joints in a fixed 
position and in cements and wall-plasters. 



Magnesium sulphate, MgS0 4 

Magnesium sulphate heptahydrate, MgS0 4 .7H a O, is the familiar 
substance, 'Epsom salt'. It occurs in springs at Epsom and Bath and 
is usually prepared from the mineral, kieserite, MgS0 4 .H 2 0, found 
at Stassfurt. It acts as a mild purgative. 

In the laboratory it may be prepared by the method described on 

p. 230. 

Zinc sulphate, ZnS0 4 

This salt is usually encountered as the heptahydrate, ZnS0 4 .7H a O, 
'white vitriol'. It can be prepared in the laboratory from zinc, zinc 
oxide or zinc carbonate, and its preparation is fully described on 

Its transparent crystals are very soluble in water (138 g in 100 g 
water at 10°C) and the salt is used as an emetic and for the treatment 
of certain skin diseases. 

IronfEl) sulphate, FeS0 4 

Iron(II) sulphate heptahydrate, FeS0 4 .7H 2 0, is known as 'green 
vitriol'. It is usually prepared in the laboratory by the action of iron 
(wire, filings or borings) on dilute sulphuric acid (p. 230). 

In industry, it is obtained by the action of air and water on the 
mineral, iron pyrites, FeS.. 

2FeS a + 70 a + 2H 2 -* 2FeS0 4 + 2H 2 S0 4 
The sulphuric acid is neutralised by scrap iron and the iron(II) sul- 
phate is crystallised. 

Fe + H 2 SO« -► FeSO« + H a 
Action of heat. On heating, iron(II) sulphate first loses its water of 
crystallisation, the original green crystals being converted into a 
dirty-yellow anhydrous solid. 

FeS0 4 .7H 2 -* FeS0 4 + 7H 2 
When more strongly heated, it gives off sulphur dioxide (test— 
decolorisation of potassium permanganate solution) in addition to 
white fumes of sulphur trioxide, and leaves a reddish-brown solid, 
iron(III) oxide, Fe a O a , 'jewellers' rouge'. This is used in pigments 
(Venetian red, red ochre) and as a polishing powder. 
2FeS0 4 -»- Fe a O a + SO s + SO„ 
Sulphuric acid was prepared by Glauber (1648) by distilling ironfll) 
sulphate crystals. The sulphur trioxide given off in the second stage 
reacted with the water driven off in the first. 

H a O + S0 3 -»-H 2 S0 4 
Iron(II) sulphate is used in the brown ring test for nitrates (p. 436) 
and it gives a similar colour with nitrogen monoxide (p. 448). 



Like all iron(TI) salts, iron(II) sulphate is a reducing agent, oper- 
ating by electron loss, which converts iron(II) ions, Fe 2+ , into 
iron(lll) ions, Fe 3+ . For example, it reduces nitric acid to nitrogen 
monoxide and chlorine to its ions. The two oxidising agents accept 
the electrons lost by the iron(H) ions. 

6Fe*+ + 6H+ + 2HNO a -»- 6Fe 3+ + 4H a O + 2NO 
2Fe 8+ + Cl a ->- 2Fe 3+ + 2CI- 

The iron(lI) sulphate is usually used in solution in dilute sulphuric 
acid to prevent hydrolysis. 

When exposed to air, iron(II) sulphate crystals become covered 
with a brownish deposit of a basic iron(TII) sulphate, by a reaction 
of the type: 

12FeS0 4 + 6H t O + 30, -► 4{Fe 2 (S0 4 ) a .Fe(OH) s } 
(from the air) 
Large quantities of iron(II) sulphate are used with gallic acid in 
the manufacture of ink. This recipe has been known for more than 
2000 years. 

Ironflll) sulphate, Fe^SO^ 

This salt may be prepared by oxidising iron(II) sulphate by nitric 
acid in the presence of sulphuric acid (equation above). 

It forms alums, for example K 2 S0 4 .Fe a (S0 1 ) 3 .24H a O, which are 
more important than iron(III) sulphate itself, because they can be 
more readily purified by crystallisation. 

Copper(II) sulphate (cupric sulphate), CuS0 4 

'Blue vitriol', CuS0 4 .5H a O, is copper(II) sulphate pentahydrate. 
The preparation of the salt from copper is fully described on p. 231 ; 
it may also be prepared from the oxide or carbonate of the metal and 
dilute sulphuric acid (p. 235). 

On the large scale, it is made by first heating scrap copper with 

Cu + S -> CuS 
and then oxidising the sulphide by heating it with access of air. 

CuS + 20 a — > CuS0 4 
The sulphate is then crystallised. 

When heated, the pentahydrate loses water of crystallisation and 
leaves white anydrous copper(II) sulphate (p. 208). 
CuS0 4 .5H a O -> CuS0 4 + 5H a O 

The formation of the blue pentahydrate from the anhydrous salt is 
used as a test for water. 

Use. Copper(n) sulphate is used in making washes such as 'Bor- 
deaux mixture' (11 parts of lime and 16 parts of copper sulphate in 



1000 parts of water), used in spraying vines and potatoes to kill 
moulds which would injure the plants. It is also used in the manu- 
facture of certain green pigments. 


The general properties of sulphites are discussed on p. 404. 

Calcium hydrogen sulphite, Ca(HSO a ) a 

This is prepared by passing sulphur dioxide into milk of lime (a 
paste of slaked lime and water) and is used for bleaching the pulp in 

Ca(OH) a + 2SO a -> Ca(HS0 3 ) a 


1. Give an account of the preparation and properties of hydrogen 

sulphide. . ...... 

If a specimen of hydrogen sulphide were contaminated with hydrogen, 
how could you obtain the hydrogen sulphide free from hydrogen? (C.) 

2. Starting from sulphur, describe how you could prepare specimens of 
(a) plastic sulphur; (b) sulphur dioxide; (c) sulphur trioxide; and (d) hydro- 
gen sulphide. (N.U.J.B.) 

3. Starting with roll sulphur, how would you prepare: 

(a) Rhombic crystals of sulphur? 

(b) Monoclinic (prismatic) crystals of sulphur? 

(c) Plastic sulphur? . . 

Mention two other elements which, like sulphur, exist in more than one 

variety. (C.) 

4 Describe one laboratory method or preparing and collecting hydrogen 
sulphide, and mention a suitable drying agent. What is the effect of passing 
hydrogen sulphide into solutions of (a) copper sulphate; (6) iron(III) 
chloride; (r) chlorine; (d) ammonia; (e) litmus? 

Briefly describe what happens when hydrogen sulphide burns in (a) excess 
of air; (b) a deficit of air. (B.) 

5 Give an account of the important properties of hydrogen sulphide. 
A specimen of this gas prepared from iron(II) sulphide is found on analysis 
to contain 10% by volume of free hydrogen. Assuming that the iron(II) 
sulphide contained no other impurity than metallic iron, calculate the 
percentage of free iron present. (L.) 

6 From what sources is sulphur obtained? How can sulphur be used 
for the preparation of (a) sulphuric acid; (b) sulphurous acid? (O and C.) 

7 How would you prepare in the laboratory hydrogen sulphide gas? 
Sketch the apparatus. What is the effect of the gas on (a) a solution of lead 
nitrate; (/>) sulphur dioxide; (c) bromine water? (N.U.J.B.) 

8. Give a short account of the chemical reactions which take place in the 
manufacture of sulphuric acid. 

Describe experiments illustrating the action of this acid on metals. (U. 
and C.) 



9. Describe the properties of sulphuric acid. 

Why is this compound regarded as (a) an acid; (A) a dibasic acid? (O. 
and C.) 

10. How would you prepare a quantity of dry sulphur dioxide? How 
may it be shosvn that the formula for this gas is SO,? (N.U.J.B.) 

11. Describe briefly two distinct methods which could be used for the 
preparation of sulphur dioxide. 

How is sulphur dioxide converted into sulphuric acid in the 'contact' 

What simple experiment shows that sulphur dioxide contains its own 
volume of oxygen? 

12. Describe the preparation and collection of dry sulphur dioxide in 
the laboratory. Mention, without giving details of the manufacturing plants 
how it is prepared on the industrial scale. What is the action of sulphur 
dioxide on (a) water; (Z>) oxygen; (c) chlorine; (d) hydrogen sulphide; 
(e) nitrogen dioxide? (B.) 

1 3. Describe the preparation and collection of sulphur dioxide. Describe 
an experiment to show that sulphur dioxide contains its own volume of 
oxygen. What additional information would you require in order to deter- 
mine the molecular formula of sulphur dioxide? Show clearly how you 
would use the results of the experiment, and the additional information in 
determining this formula. (N.U.J.B.) 

14. Describe the reaction which takes place when copper is heated with 
concentrated sulphuric acid. The resulting gas is passed into (a) litmus 
solution; (6) chlorine water; (c) a solution of hydrogen sulphide. What 
would be observed in each case and what explanations would you give of 
the results obtained? (L.) 

15. Describe fully how to prepare and collect in the laboratory sulphur 
dioxide from sulphuric acid. How would you show the action of this gas 
on (a) chlorine water; (6) moist hydrogen sulphide? 

What takes place when a solution of the gas is allowed to stand in con- 
tact with air? 
Explain the above reactions by equations or otherwise. (N.U.J.B.) 

16. What is meant by the term allotropy? Describe the preparation of 
two allotropic forms of sulphur. 

Starting from sulphur, how would you obtain fairly pure samples of 
(a) sulphur dioxide; (b) sulphur trioxide? (C.W.B.) 

17. Describe, with a diagram, how you would prepare and collect 
hydrogen sulphide in the laboratory. 

Describe how hydrogen sulphide reacts with (a) sulphur dioxide; 
(Z>) iron(ril) chloride solution. 

When electric sparks from an induction coil are passed for some time 
through a volume of hydrogen sulphide the gas is decomposed, sulphur is 
deposited on the sides of the vessel and on cooling to the original con- 
ditions hydrogen remains, the volume of which is equal to that of the 
hydrogen sulphide. The vapour density of the hydrogen sulphide being 17 
calculate from these facts the formula of hydrogen sulphide. (H, 1 ; S, 32.) 

18. Explain the construction and the working of a Kipp's apparatus for 
generating hydrogen sulphide. Describe and explain the effect of passing 
the gas through aqueous solutions of (a) copper sulphate; (A) blue litmus; 
(c) chlorine. What happens if the resulting solution from (6) is boiled ? (L.) 

Chapter 32 

Nitrogen and Ammonia 

(For nitrogen in periodicity, see p. 84) 



ABOUT four-fifths of the atmosphere is free nitrogen. The element 
also occurs combined in the form of sodium nitrate, Chile salt- 
petre, NaN0 3 , as a mineral deposit in Chile, and distributed every- 
where in the soil in minute quantities as ammonium sulphate 
(NHJaSO^ and sodium nitrate, NaN0 3 , potassium nitrate, KN0 3 , 
and calcium nitrate Ca(N0 3 ) 2 . (The very great importance of these 
compounds of nitrogen in maintaining the fertility of the soil is dis- 
cussed on p. 147.) 

Combined nitrogen is always found as a constituent of the living 
matter of plants and animals. 

Preparation of nitrogen from the atmosphere 

The gases present in dry air are oxygen, about 21% by volume, 
carbon dioxide, about 0.03% by volume, and 'atmospheric nitrogen', 
about 79% by volume. The first two of these gases can be removed 
and the nitrogen collected by the apparatus of Fig. 135. 

Copper to absorb 
I Oxygen 


\Caustic soda solution 
to absorb Carbon Dioxide. 

Fig. 135. 

Preparation of nitrogen from tho air. 




Na 2 C0 3 + H,0 

The equations arc: 

Absorption of carbon dioxide 2NaOH + CO* 
Absorption of oxygen 2Cu + 2 

If the nitrogen is required dry, it may, after leaving the heated 
copper, be passed through a U-tube containing glass beads wetted 
with concentrated sulphuric acid to dry it and then collected over 

The product of this experiment is not pure nitrogen. It contains 
about 1 % by volume of the 'rare gases', chiefly argon, the removal of 
which is not possible by chemical methods. The presence of these 
gases makes 'atmospheric nitrogen' denser than the pure gas. 

Another method of preparing 'atmospheric' nitrogen is to absorb 
both carbon dioxide and oxygen together by shaking air with a 
solution of pyrogallol in sodium hydroxide solution. The sodium 
hydroxide absorbs the carbon dioxide and the pyrogallol absorbs 
the oxygen to form an oxidation product of itself. 

This method is, however, only suitable for the preparation of small 
samples of 'atmospheric nitrogen'. 

Preparation of nitrogen by heating ammonium nitrite 

A solution of ammonium nitrite readily decomposes on slight 
warming to give nitrogen. This decomposition occurs slowly' at 
ordinary temperatures, so that neither ammonium nitrite itself, nor 

Solution of 
Sodium Nitrite and 
Ammonium Chloride 


Fio. 136. 
Preparation of nitrogen by a chemical method. 

its solution in water, should be kept in stock. The compound is pre- 
pared as required by a double decomposition reaction between 
sodium nitrite and ammonium chloride. 

NaNO, + NH 4 C1 
69 g 53.5 g 

NaCl + NH 4 N0 2 




Weigh out the two compounds in these proportions. 14 g of sodium 
nitrite and 1 1 g of ammonium chloride will be suitable weights. 
Place the compounds in a round flask, add 350 cm 3 of water, fit up 
the apparatus as in Fig. 136 and heat gently. As the solution becomes 
warm, rapid effervescence occurs and the nitrogen evolved may be 
collected over water. 

NH 4 N0 2 (aq) -»- N„(g) + 2H 2 0(1) 

Other chemical methods of preparation or nitrogen 
The action of chlorine on ammonia. 

3C1, + 8NH 3 -*■ N g + 6NH 4 C1 

Passing ammonia gas over heated copper(II) oxide (see p. 429). 

2NH a + 3CuO-> 3Cu + N 2 + 3H 2 
Reduction of oxides of nitrogen by heated copper, e.g., 

2Cu(c) + 2NO(g) -»■ 2CuO(c) + N,(g) 
These methods are all much less convenient than the heating of 
ammonium nitrite. 

Tests for nitrogen 

At ordinary temperatures, nitrogen is so inert that no positive 
tests can be applied. We can only show a given gas to be nitrogen by 
elimination of other possibilities. 

Lighted splint. Place a lighted splint into a gas-jar of the gas. It is 
extinguished and the gas does not burn. It cannot, therefore, be any 
gas which supports combustion, e.g., oxygen, dinitrogen oxide, or any 
combustible gas, e.g., hydrogen sulphide, carbon monoxide, hydrogen 

Smelt. The gas has no smell. This distinguishes it from gases such 
as sulphur dioxide, ammonia, hydrochloric acid gas. 

Action of lime-water. After the above tests the only gas with which 
nitrogen may be confused is carbon dioxide. To distinguish it from 
this, add lime-water and shake. Nitrogen leaves the lime-water un- 
changed; with carbon dioxide, the lime-water is turned milky. 

Properties of nitrogen 

Nitrogen is colourless and odourless. It is slightly lighter than air 
and only slightly soluble in water (about 2 volumes of the gas 
dissolve in 100 volumes of water at ordinary temperature). 

Under ordinary conditions the gas is very inert, but, by applying 
the results of much research, it has been made to combine with 
hydrogen to produce ammonia. 

N 2 + 3H 2 ^2NH 3 
For this reaction applied in Haber's Process, see p. 431. 





Nitrogen will combine directly with many metals forming nitrides 

3Mg + N g ->Mg 3 N )1 

To illustrate this, burn some magnesium ribbon in a crucible and 
allow the product to cool. Add a few drops of water and smell the 
mixture. The choking smell is that of ammonia. It is evolved by the 
action of water on the magnesium nitride which was formed, in small 
amount, by combination of the magnesium with nitrogen of the air. 

Mg 3 N, + 6H s O 
magnesium water 

->-2NH 3 +3Mg(OH) 2 
ammonia magnesium 


This hydride of nitrogen, NH 3 , can be made in very small amounts 
by heating nitrogenous organic materials such as hoofs and horns of 
animals. Its old name was, in fact, 'spirit of hartshorn'. 

Preparation of ammonia 

Ammonia may be prepared in the laboratory by heating any 
ammonium salt with an alkali. Usually a mixture of ammonium 
chloride and calcium hydroxide (slaked lime, the cheapest alkali) is 
used. Both are solids so they must be thoroughly ground first to give 
a very fine mixture in which the reaction can occur satisfactorily. 

Ca(OH) i! + 2NH 4 C1 -> CaCl a + 2H 8 + 2NH,(g) 
74 g 2 x 53.5 g 

107 g 

An excess of the slaked lime is preferable. Weigh out 25 g of 
slaked lime and 16 g of ammonium chloride. Grind the mixture well 
in a mortar, place it in a round flask of resistance glass and set up 
apparatus as in Fig. 137. The neck of the flask should slope towards 
A as shown, because water will condense and, if allowed to run back 
on to the hot flask, might break it. Heat the flask. Ammonia gas is 
evolved. It is dried by a rather unusual drying agent, quicklime, CaO, 
because it reacts with all the usual drying agents. Concentrated 
sulphuric acid is acidic and would absorb the gas, forming a salt, e.g., 

2NH 3 + H 2 S0 4 -> (NHJ.SO, 

while it reacts with calcium chloride, forming solid complex com- 
pounds, e.g., 

CaCI 2 + 4NH 3 — > CaCI 2 .4NH 3 



calcium hydroxide and 
ammonium chloride 


Card cover 


quicklime to 
dry the 

Fig. 137. 
Preparation of ammonia gas. 

Ammonia is less dense than air and very soluble in water, so it is 
collected as shown by upward delivery. 

Instead of calcium hydroxide, sodium hydroxide (or potassium 
hydroxide) solution may be used, in which case the flask would be 
placed in the vertical position and heated on a tripod and gauze. 

NaOH + NH 4 C1 -»- NH 8 + H,0 + NaCl 
or OH- + NH 4 + -*■ NH 3 + H 2 

Ammonium sulphate may be used instead of ammonium chloride. 

Ca(OH) a + (NH 4 ) s S0 4 -»- CaS0 4 + 2H 2 + 2NH 3 
2NaOH + (NH 4 ) 8 S0 4 -* Na,S0 4 + 2H 2 + 2NH S 

To obtain ammonia gas in quantity 

If several gas-jars of ammonia are required, it is very convenient to 
fill these by heating a concentrated solution of the gas in water. In 
this case the flask containing the ammonium chloride and slaked lime 
in the previous experiment is replaced by a vertical flask containing a 
concentrated solution of ammonia. N.B. This is not, strictly speaking, 
a preparation of ammonia gas but merely obtaining the gas from its 
solution in water. 





Tests for ammonia 

Smell. A characteristic choking smell. The choking smell 
is due to the fact that the gas temporarily paralyses the 
respiratory muscles and breathing is checked. In large 
quantities, the gas causes asphyxiation. 

Action with litmus. Expose damp red litmus paper to the 
gas. It is turned blue. Ammonia is the only common alkaline 

Properties of ammonia 

Appearance. A colourless gas. 
Density. Lighter than air. 

Density relative to hydrogen, 8.5. 
Density of air relative to hydrogen, 14.4. 

Solubility of ammonia in water. The fountain experiment 

The very great solubility of ammonia in water is illustrated in the 
fountain experiment (Fig. 138). Replace the gas-jar of Fig. 137 by a 
dry, thick-walled flask of about 1500 cm 3 capacity and pass ammonia 

into it for some time. (It is better 
to supply another flask with fresh 
reaction mixture, to give off a satis- 
factory stream of ammonia.) Fit 
the flask with a rubber stopper 
carrying tubes and clips as shown. 
Place the tubes and clips under 
water, open clip B for a moment, 
close it and allow the few drops of 
water which have entered to run 
down into the round part of the 
flask. Then replace the tubes and 
clips under water and open clip A. 
A fountain will at once play, as in 
sketch, and will continue until the 
flask is as full of water as it was 
formerly full of ammonia. 
The alkaline nature of ammonia can be shown in this experiment 
by adding a little litmus solution to the water in the trough and 
making it turn red by the addition of a drop of acid. When the 
ammonia dissolves in the litmus solution, the latter is turned blue. 
Explanation. Ammonia has the highest solubility of all known gases 
(about 800 vol. of gas in 1 vol. of water at 15°C). The first few drops 
of water, which entered when clip B was opened, dissolved nearly 

Fio. 138. 
Fountain experiment. 

800 times their own volume of ammonia. This reduced the gas pres- 
sure inside the flask to only a fraction of its former value, atmospheric 
pressure. As soon as the clip A was opened, the water was forced 
into the flask because the atmospheric pressure from outside over- 
came the resistance of the reduced gas pressure inside the flask. The 
water, entering as a fountain, dissolved the remaining ammonia, 
maintaining the fountain until only air was left in the flask. (A thin- 
walled flat-bottomed flask must not be used for this experiment; the 
reduction of pressure inside would almost certainly cause it to collapse 

Action of ammonia with hydrogen chloride 

Place a gas-jar of ammonia over a gas-jar of hydrogen chloride, 
remove the covers and allow the gases to mix. Dense white fumes will 
be seen which will settle to a white solid, sal-ammoniac or ammonium 
chloride, on the sides of the gas-jar. 

NH 3 (g) + HCl(g)-*NH 4 CI(c) 

Ammonia as a reducing agent 

Action of ammonia on chlorine. This is fully considered on p. 1 14. 

Action of ammonia on heated copper(ll) oxide. Set up the apparatus 

as shown in Fig. 139. The colourless liquid collecting at A is water 


Granular copper ( II ) oxide nitrogen 


Gas Furnace 

water water 

Fio. 139. 
Action of ammonia on copperflO oxide. 

(for tests, see p. 270) and the colourless gas at B is nitrogen (for tests 
see p. 425). The ammonia has reduced the copper(II) oxide to copper 
and has itself been oxidised to nitrogen and water. 

3CuO(c) -I- 2NH 3 (g) -> 3Cu(c) + 3H g O + N 2 (g) 
Combustion of ammonia. Ammonia will burn in an atmosphere of 
air slightly enriched by oxygen but not in air alone. The chief pro- 
ducts are nitrogen and water. 

4NH 3 + 30 s -> 2N 2 + 6H 2 Q 



Liquefaction of ammonia gas 

Ammonia gas can be liquefied at ordinary temperatures by com- 
pression. The colourless liquid boils at about — 33°C under ordinary 
atmospheric pressures. 

Formula of ammonia gas — see p. 1 14 and p. 74. 

Ammonium hydroxide, NH 4 OH 

Set up apparatus as in Fig. 140, and heat the flask gently as in the 
preparation of ammonia gas. The rim of the inverted funnel should 
just touch the surface of the water. This is a device to prevent the 
water from 'sucking back' into the flask. After a time the water in 
the beaker will be found to have acquired the smell of ammonia gas 
which has dissolved in it. The solution is known as ammonium hydrox- 
ide. At 0°C and ordinary pressure, one volume of water dissolves 

ammonium chloride and 
calcium hydroxide 


Fig. 140. 
Preparation of ammonium hydroxide solution. 

about 1000 volumes of the gas. Some of this gas combines chemically 
with the water to form ammonium hydroxide, which acts as a weak 
base, ionising slightly (about 0.4% in normal solution), as: 

NH 3 + H 2 v± NH 4 + + OH- 
The hydroxyl ion, OH-, gives the solution its alkaline reaction to- 
wards litmus and many properties resembling those of the caustic 
alkalis, Na + OH- and K+OH". Like them, it will precipitate in- 
soluble metallic hydroxides when mixed with solutions of salts or the 
metals, e.g., 

3K.OH + FeCls ->■ Fe(OH) 3 1 

3NH 4 OH + FeCl 3 -»- Fe(OH) 3 , 

+ 3KC1 
+ 3NH 4 C1 



It will also neutralise acids forming ammonium salts, which can be 

crystallised out and are generally similar to ordinary metallic salts, e.g., 

NaOH + HCl-J-NaCl + HjO \ 

NH 4 OH + HC1 -»- NH 4 C1 + H,0/ 

2NaOH + H 2 S0 4 -*■ Na 2 S0 4 + 2H a O \ 

2NH 4 OH + H 2 S0 4 -> (NH«) a S0 4 + 2H 2 Oj 

Notice that ammonium salts, though often written in molecular form, 

are actually electrovalent compounds, containing the ammonium ion, 

NH 4 + , in combination with a corresponding amount of an acidic 

ion, such as CI" or NO a _ . Since the ammonia molecule NH S can 

combine with a hydrogen ion (proton), it must be regarded as a base. 

NH 3 + H + ^NH 4 + (Seep. 70) 
It is a much weaker base than the hydroxide ion so ammonium salts 
of strong acids are considerably hydrolysed in their aqueous solu- 
tions and these solutions show appreciable acidity. 

Uses of ammonia 

(1) Ammonia solution is used in laundry work. It removes tem- 
porary hardness by precipitating the calcium ion of calcium hydrogen 
carbonate as chalk. 

Ca»+ + 2HCO„- + 20H- -*- CaCO s + 2H 2 + C0 3 »- 
It also dissolves out acids left by evaporation of perspiration from 

(2) Ammonia is converted to nitric acid. (See p. 452.) 

(3) Ammonium sulphate is made from ammonia to be used as a 
nitrogenous fertiliser. This may be done by direct neutralisation with 
sulphuric acid or by passing carbon dioxide and ammonia through a 
steam-heated mixture of calcium sulphate (anhydrite) and water. 

CaS0 4 + 2NH 3 + CO a + H 2 -* CaC0 3 + (NHJgSO, 
Chalk is filtered off and ammonium sulphate can be crystallised. 

(4) Liquid ammonia is used in refrigerating plant. 

Manufacture of ammonia. Haber's process 

This process affords a good example of the effects of various 
factors (e.g., temperature, pressure, concentration) on chemical re- 
action and the way in which industrial practice is influenced. The 
general situation is summarised in Le Chatelier's Principle, which 

If a chemical system is in equilibrium and one of the 
factors involved in the equilibrium is altered, the equi- 
librium will shift so as to tend to annul the effect of the 






Haber's process uses the reaction: 
N,(g) + 3H 2 (g) 

1 vol. 

3 vol. 

2NH 3 (g) 
2 vol. at constant T. and P. 

The reaction (from left to right) is exothermic, i.e., liberates heat. 

Effect of pressure. Ammonia is produced from its elements with 
reduction of volume. Therefore, if the system is in equilibrium and 
the pressure is then raised, the equilibrium must shift so as to tend to 
lower the pressure (Le Chatelier's Principle). To do this, the volume 
must be reduced by production of more ammonia. That is, high 
pressure favours production of ammonia. 

Effect of temperature. The formation of ammonia from its elements 
is an exothermic change. If the system is in equilibrium and the 
temperature is then lowered, the equilibrium must shift so as to tend 
to raise the temperature again (Le Chatelier's Principle). That is, heat 
must be liberated by the production of more ammonia. That is, low 
temperature favours the production of ammonia. But lowering of 
temperature reduces the rate of reaction, so it is necessary to introduce 
a catalyst which will give a sufficient reaction rate in spite of a rela- 
tively low temperature. 

Effect of concentration. If the system is in equilibrium and more 
nitrogen is then added to increase its concentration, Le Chatelier's 
Principle requires the equilibrium to shift so as to tend to reduce the 
nitrogen concentration. That is, more ammonia will be produced to 
use up nitrogen. This increases the yield of ammonia relative to 
hydrogen, and vice versa if the hydrogen concentration is increased. 
However, in practice, there is no particular advantage in using excess 
of either material; the gases are used in the theoretical proportion of 
nitrogen to hydrogen, 1 : 3 by volume. 

The conditions chosen in accordance with the above discussion 

very high pressure (200-500 atm) 
temperature about 450°C. 
catalyst: finely divided reduced iron, usually 
'promoted' by alumina. 

Hydrogen is manufactured from hydrocarbons and nitrogen from the 
air. They are mixed in the required 3 : 1 proportion by volume and 
dried (e.g., by silica gel). They are pre-heated by gases leaving the 
catalyst chamber and are then passed over the catalyst at 450°C. The 
ammonia produced is absorbed in water or liquefied by refrigeration 
and the unused gases are recirculated. 

N t + 3H 2 -»-2NH3 
The ammonia and sulphuric acid industries on the Tees have the 
advantage of immediately adjacent deposits of anhydrite (CaSO^), of 

coal from Durham, and limestone from the Tees and Wear valleys. 
There is also easy transport by sea for fertilisers to Europe. 

Action of alkalis on ammonium salts 

Any ammonium salt, if heated with sodium hydroxide (or po- 
tassium hydroxide) solution or with calcium hydroxide and a little 
water, gives off ammonia gas (which turns red litmus paper blue). 
This distinguishes ammonium salts from those of any metal. 
NH« + + OH- ->- NH 3 + H a O 

Action of heat on ammonium salts 

Sublimation. Ammonium salts are always decomposed by heat 
and, sometimes, sublime. The best example of sublimation is provided 
by ammonium chloride. 

Place a little ammonium chloride in a dry test-tube and heat gently. 
The effects shown in Fig. 141 will be observed. 

Sublimate of 


Fig. 141. 

Action of heat on ammonium chloride. 

Usually when a vapour is cooled, it condenses first to a liquid, and, 
later, on further cooling, solidifies, e.g., steam -> water -* ice. 

The characteristic feature of sublimation is that, on cooling, the 
vapour condenses directly to the solid without the intermediate 
liquid state. Usually, the converse is also true, that the subliming 
solid is converted directly to vapour on heating without an inter- 
mediate liquid stage, but in some cases (for example, that of iodine) 
melting may occur. 

Very few substances sublime. Among those which sublime are a 
few ammonium salts (especially the chloride), iodine and naphtha- 
lene. Sublimation is a very effective means of purifying them, because 



their impurities are very unlikely to sublime. The white sublimate of 
ammonium chloride in Fig. 141 will be a purer sample of the com- 
pound than the original material used for the experiment. For 
thermal dissociation, see p. 159. 

Decomposition of ammonium salts by heat. Ammonium salts of 
acids having a high proportion of oxygen are usually decomposed by 
heat, e.g. 

(i) Ammonium nitrite. 

NH 4 NO s ^-N 2 + 2H t O 
This reaction is fully dealt with on p. 424. 
(ii) Ammonium nitrate. 

NH 4 N0 8 -»-N,0 + 2H ! 

This reaction is fully considered on p. 445. 

Questions on this chapter will be found on page 453. 

Chapter 33 

Nitric Acid and the Nitrates 


THE old name for nitric acid was 'aqua fortis'— strong water. It 
was so called because it attacks so many substances, including 
almost all the metals. 

Preparation of nitric acid 

Set up the apparatus of Fig. 142. Into the bulb of the retort put 
some potassium nitrate crystals and add concentrated sulphuric acid. 

Fio. 142. 
Preparation of nitric acid. 

Heat the retort gently. The potassium nitrate gradually dissolves and 
effervescence occurs. 

KNO a + H a S0 4 


K.HSO4 + HNO s 



The nitric acid distils and collects in the cooled receiver as a yellow 
liquid, while drops of the acid can be seen running down the bulb and 
neck of the retort. The brown fumes are nitrogen dioxide formed by 
slight decomposition of the nitric acid by heat, 

4HN0 3 -»■ 2H 2 + 4NO, + O, 

and they impart a yellow colour to the acid by dissolving in it. 

This reaction is a general one. Any metallic nitrate, when heated 
with concentrated sulphuric acid, gives off nitric acid, e.g., 

NaN0 3 + H,SO« -> NaHS0 4 + HNO s 
sodium nitrate 

Pb(NO s ) g + H,S0 4 -> PbSO« + 2HNO s 
lead(ll) nitrate 


Some nitric acid is manufactured by the above laboratory process. 
Iron retorts are used, because this metal is only slightly attacked by 
the acid, which is condensed in silica condensers. The nitrate em- 
ployed is sodium nitrate, NaN0 3 , 'Chile saltpetre* (see p. 442). An- 
other method of manufacture is oxidation of ammonia (p. 452). 
Nitric acid is used mainly for the manufacture of explosives and dyes. 

Test for soluble nitrates 

Crush a few potassium nitrate crystals in a mortar and put them 
into a test-tube and add water to a depth of about 2 cm. Shake to 
dissolve the potassium nitrate. Add a little dilute sulphuric acid and 
then two or three crystals of iron(II) sulphate, which have also been 
crushed. Shake to dissolve them. Hold the test-tube in a slanting 
position and pour a slow continuous stream of concentrated sul- 
phuric acid down the side. (Care!) It will form a separate layer under- 
neath the aqueous layer and, at the junction of the two, a brown ring 
will be seen. This brown ring is the characteristic test for a soluble 

Explanation. The concentrated sulphuric acid and the nitrate yield 
nitric acid. 

KNO s -f H 2 S0 4 -»- KHS0 4 + HNO s 
The nitric acid is then reduced by some of the iron(II) sulphate to 
nitrogen monoxide, NO. 

6FeSO« + 2HNO-, + 3H,SO« -► 3Fe 2 (S0 4 ) 3 + 4H 2 + 2NO 

iron(lll) sulphate 
The nitric oxide reacts with more ironfjl) sulphate to give the brown 
compound FeS0 4 .NO, which appears as the brown ring. 
FeS0 4 + NO ->- FeS0 4 .NO 



Properties of nitric acid 

Nitric acid is a colourless, fuming liquid of density 1.5 g/cm* and 
boiling point 85°C at ordinary atmospheric pressure. The ordinary 
concentrated nitric acid, as sold, contains about 70% by weight of the 
pure acid and 30% of water. The pure acid is corrosive and destroys 
organic matter very readily. The skin is stained yellow by it and, if the 
acid is left in contact with it for even a very short time, the skin is 

Chemical properties of nitric acid. Nitric acid can behave 
chemically in two ways: 

It is (1) a very strong acid, 

(2) a powerful oxidising agent. 

Nitric acid acting as an acid 

Nitric acid is a very strong acid, being almost completely ionised in 
dilute solution with the production of hydrogen ion and the nitrate 

ion, N0 3 _ . 

HNO s ^ H + + N0 3 ~ 

This ionisation confers on it the usual acidic properties, modified to 
some extent by the powerful oxidising action of the acid. 

(a) Nitric acid neutralises bases, forming metallic nitrates, e.g., 

K+OH- + H + N0 4 - -* K + NO s - + H 2 
Cu 2+ O z - + 2(H + N0 3 -) -*■ Cu* + (NO,-), + H 8 

A full account of reactions of this type is given in Chapter 16, pp. 

(b) Nitric acid liberates carbon dioxide in reaction with metallic 
carbonate, as: 

C0 3 *" + 2H + -»- H 2 + CO, 
It is useful to remember that all metallic nitrates are soluble in water. 

(c) It is characteristic of strong acids that, when dilute, they react 
with the more electropositive metals, liberating hydrogen, as: 

Zn (or Mg) + 2H + -»- Zn s+ (or Mg J+ ) + H, 
This reaction cannot occur in this simple form with nitric acid; it is 
complicated by the fact that nitric acid is a powerful oxidising agent. 
Any hydrogen initially produced by the action of a metal on nitric 
acid is at once oxidised by more of the acid to water, and the reduc- 
tion products of the acid are liberated. These may include nitrogen 
dioxide, nitrogen monoxide or even ammonia, which is at once con- 
verted, by excess of the acid, to ammonium nitrate. Exact equations 
cannot be written because the reactions are always complex. If, 



however, copper is used with the concentrated acid, the principal 
reaction is: 

Cu + 4HNO, -»• Cu(N0 3 ) 2 + 2H 2 + 2NO, 

If the concentrated acid is diluted with its own volume of water, 
copper gives the principal reaction: 

3Cu + 8HN0 3 ->- 3Cu(N03), + 4H s O + 2NO 

This is the usual laboratory preparation of nitrogen monoxide. The 
important point is that, if nitric acid acts with a metal under ordinary ex- 
perimental conditions, the product is never hydrogen; it is a reduction 
product of the acid, such as nitrogen dioxide, nitrogen monoxide 
or ammonium nitrate. If, however, very dilute nitric acid is used 
(about 1%) with magnesium or manganese, some hydrogen will be 
produced, escaping oxidation because of the very dilute condition of 
the acid. The hydrogen is not pure, being accompanied to some 
extent by reduction products of the acid. 

Nitric acid as an oxidising agent 

Like oxidising agents in general, nitric acid acts as an acceptor of 
electrons. It can do this in several different ways. Two of the more 
important ones are: 

4HN0 3 + 2e~ -* 2NO a " + 2H a O + 2NO a 
8HNO3 + 6e- -> 6NO3- + 4H a O + 2NO 

giving nitrogen dioxide and nitrogen monoxide. The electrons are 
supplied by the reducing agent which takes part in the reaction. This 
is often a metal, e.g., copper. This forms ions, as: Cu — »■ Cu 2+ + 2e~. 
By combining this ionisation with the two equations given above, the 
reactions shown between copper and nitric acid in the last paragraph 
are obtained. Other metals give similar behaviour, varying in detail 
with the nature of the metal, the concentration of the acid and the 
temperature employed. 

Nitric acid also oxidises certain non-metallic elements and certain 
compounds; examples are given in the following paragraphs. 

The oxidising action of nitric acid on sulphur 

If concentrated nitric acid (which contains 70% HN0 3 and 30% 
water), is warmed with powdered sulphur, brown fumes of nitrogen 
dioxide are evolved and sulphuric acid is left in solution. Bromine 
catalyses this reaction. 

S + 6HN0 3 ->- H 2 SO« + 2H 2 + 6NO, 

In electronic terms, the presence of the powerful oxidising agent, hot 
concentrated nitric acid, enables the sulphur-water system to make 
electrons available and so act as a reducing agent. 

S + 4H,0 -v H s S0 4 -f 6H + -I- 6e~ 


The electrons arc accepted by the acid in its capacity of oxidising 

6HN0 3 +- 6H + + 6e~ -»- 6H t O 4- 6NO, 
The molecular equation above is given by adding these two electronic 

The oxidising action of nitric acid on an iron(II) salt 

Dissolve a few crystals of iron(II) sulphate in dilute sulphuric acid 
in a test-tube. Add a little concentrated nitric acid and heat. 1 Brown 
fumes of nitrogen dioxide are seen and a brown or yellow solution 
is left, instead of the original pale green solution. 

The nitric acid has oxidised the green iron(II) sulphate to brown 
or yellow iron(III) sulphate and has itself been reduced to nitrogen 
monoxide which, in the air, forms nitrogen dioxide (p. 448). 

6FeS0 4 + 3H.S0 4 + 2HN0 3 -> 3Fe 2 (SO«) 3 + 4H 2 + 2NO 
2NO + 2 -v 2NO s 
of the 

An important point to notice about nitric acid as an 
oxidiser is that it introduces no solid into a mixture. The 
acid itself is volatile and its reduction products, oxides of 
nitrogen, are gaseous, while the other product, water, can 
also be evaporated off. Thus, it leaves no solids to compli- 
cate purification of the product of oxidation. 

Other oxidising actions of hot, concentrated nitric acid are: 
Carbon to carbon dioxide. 

C + 4HN0 3 ->- 2H 2 + 4NO, + CO, 
Red phosphorus to orthophosphoric acid. 

P + 5HN0 3 -> H 3 P0 4 + 5N0 2 + H t O 



The action of heat on concentrated nitric acid 

A clay pipe is used for this experiment because a fairly high 
temperature is needed and the liquid nitric acid would break a heated 
glass tube. Fit up apparatus as in Fig. 143, omitting, for the present, 
the boiling-tube, and raising the end of the pipe stem above the water 

1 Before heating, the solution will be very dark brown or black. For an ex- 
planation of this, sec the 'brown ring' test (p. 436). 



Nitric acid 

Fig. 143. 
Decomposition of nitric acid by heat. 

Heat the pipe stem to red heat at one point with a Bunsen burner 
and pour concentrated nitric acid into the bowl of the pipe. Clouds 
of brown fumes emerging from the pipe stem show that nitrogen 
dioxide is produced. 

Now, complete the apparatus as in Fig. 143, making sure that the 
end of the pipe stem is the least possible distance below the water-level 
and starting with the boiling-tube full of water. As the nitric acid 
runs over the red-hot pipe stem, a colourless gas collects in the boiling- 
tube. Test it with a glowing splint. It is rekindled. The gas is oxygen. 
(Nitrogen dioxide, also produced as above, is soluble in water.) 

The third product of the decomposition of the acid is water. 

4HN0 3 -»• 2H 2 + 4NO, + O, 













Nitrates of these metals are decomposed by 
heat to the nitrite and oxygen. 

Nitrates of these metals are decomposed on 
heating to the oxide of the metal, nitrogen 
dioxide and oxygen. 

Nitrates of these metals are decomposed on 
heating to the metal, nitrogen dioxide and 




Aluminium nitrate, A1(N0 3 ) 3 .9H 2 

This is an unimportant salt. It can be prepared by dissolving 
freshly prepared aluminium hydroxide in dilute nitric acid and 
crystallising in the usual way. 

3HN0 3 -* Al(NO s ) 3 + 3H,0 


Al(OH) 3 

It shows the usual behaviour of the nitrate of a heavy metal when 

4A1(N0 3 ) 3 -+ 2A1,0, + 12NO, + 30» 

Ammonium nitrate, N 1 1 ..\0 :; 

This compound may be made in the laboratory by neutralisation 
of ammonium hydroxide by nitric acid (p. 233). 

NH 4 OH + HNO a -»- NH,N0 3 |- H a O 
It is colourless and very soluble in water. When it dissolves, heat is 
absorbed and, by dissolving a large quantity of the salt in water, a 
liquid of low temperature is rapidly obtained and may be used as a 

Ammonium nitrate is decomposed by heat into dinitrogen oxide 
and water (p. 445). 

NH 4 N0 3 -> N t O + 2H,0 

If the experiment is carried out in a test-tube, the usual test for 
dinitrogen oxide (rekindling of a glowing splint) will be masked by 
the steam which is also given off. If the heating is carried on to de- 
compose all the ammonium nitrate, there will be no residue and the 
last traces of the salt will decompose with explosion. 

A mixture of ammonium nitrate and aluminium powder is used as 
an explosive, 'ammonal'. When it is detonated, the following re- 
action occurs. 

2A1 + 3NH<N0 3 -*- A1,0 3 + 3N 8 + 6H,0 
The gaseous nitrogen and steam, having a volume many times greater 
than that of the original solids, produce a very high pressure and 
hence an explosion. 

Potassium nitrate (nitre, saltpetre), KNO a 

In the laboratory, this salt may be obtained by neutralisation 
(p. 233). 

KOH !- HN0 3 ->* KNO s + H t O 

On the industrial scale, it is prepared by double decomposition 
between potassium chloride (from the Stassfurt deposits) and sodium 
nitrate (from Chile). 

KC1 + NaN0 3 -*■ KN0 3 + NaCl 
Boiling saturated solutions of potassium chloride and sodium nitrate 
are used and sodium chloride, being the least soluble of the four salts 
at this temperature, crystallises and is filtered off. On cooling to 



ordinary temperature, potassium nitrate crystallises as it is the least 
soluble of the four salts at this temperature. 

When heated, potassium nitrate melts to a colourless liquid and 
decomposes slowly, liberating oxygen (test: glowing splint rekindled) 
and leaving, when cool, a pale yellow solid, potassium nitrite. 
2KN0 3 -► 2KNO a + O a 

Potassium nitrate is chiefly used for the making of fireworks and 
gunpowder, which usually contains about one part of charcoal, one 
part of sulphur and six parts of potassium nitrate (by weight). When 
ignited, the mixture burns rapidly, producing nitrogen, oxides of 
carbon and sulphur and other gases. These hot gases occupy a much 
greater volume than the original solids, and a very great pressure is 
set up which is used for propulsion or disruption. 

Sodium nitrate (Chile saltpetre), NaN0 3 

Large deposits of sodium nitrate, mixed with clay, occur in Chile, 
as 'caliche'. The material is broken up by blasting. The sodium 
nitrate is extracted by dissolving it out in water and evaporating the 
solutions by the heat of the sun. The area where the nitrates are 
found is practically rainless and water is supplied for the process by 

In the laboratory, sodium nitrate may be prepared by neutralising 
caustic soda with nitric acid (p. 233). 

NaOH + HN0 3 -> NaN0 3 + H a O 

When heated, sodium nitrate behaves exactly like potassium 
nitrate (see last section). 

2NaN0 3 ->- 2NaNO a + 3 

The most important use of sodium nitrate is its application to the 
land as a fertiliser (see p. 452). It is rapid in action and is applied in 

The Chilean deposits used to be the only source of nitrogenous 
'chemical' fertiliser, but the Haber process (p. 431) now furnishes 
alternative supplies. 

Large quantities of sodium nitrate are also used for the manu- 
facture of nitric acid (p. 436), potassium nitrate (see p. 441), and 
sodium nitrite. This substance is used extensively in the manu- 
facture of aniline dyes, and is made by heating sodium nitrate with 
carbon or lead. The presence of these reducing agents hastens the 
conversion of the nitrate to the nitrite, a process which is slow if heat 
alone is employed. 

NaNOa + Pb-> NaNO a + PbO 
2NaN0 3 + C -► 2NaN0 2 + CO a 



Calcium nitrate, Ca(NO s ) a 

This compound may be made in the laboratory by the action of 
nitric acid upon slaked lime or chalk. 

Ca(OH) a + 2HN0 3 
CaC0 3 + 2HN0 3 

Ca(N0 3 ) a + 2H a O 
CaiW)^ + H a O + CO, 

It is usually met with as the tetrahydrate, Ca(N0 3 ) 3 .4H a O, which 
forms white deliquescent crystals. 

In industry, some calcium nitrate (mixed with lime) is used as a 
fertiliser, 'air saltpetre'. It is prepared from the nitric acid obtained 
by oxidising ammonia. 

For the action of heat on calcium nitrate, see below. 

Nitrates of magnesium, zinc, lead, copper 

These compounds are all prepared, as explained in Chapter 16, by 
the action of nitric acid on the metal or on its oxide, hydroxide or 
carbonate. Except leadfll) nitrate, which crystallises anhydrous, 
they all form hydrated crystals. All the crystals are white in colour 
except those of copper(II) nitrate, which are blue. 

The most important property of these nitrates is the reaction they 
undergo when heated. This is as follows (taking for example lead 

Place a little lead(ll) nitrate in a test-tube and heat it. A series of 
sharp cracking sounds, almost small explosions, is heard. (This effect 
is called decrepitation.) Later, the lead(II) nitrate melts and effer- 
vesces, giving off brown fumes (nitrogen dioxide). Test the gas with 
a glowing splint. The splint is rekindled. Hence oxygen is also present 
The solid residue is reddish brown when hot and yellow when cold. 
It is litharge (lead(ll) oxide), and will probably be found to be fused 
into the glass. 

2Pb(NO a ) a -> 2PbO + 4NO a + O a 

This reaction is typical of the action of heat on the nitrates of com- 
mon heavy metals. (Note that it is also exactly analogous to the 
action of heat on nitric acid, p. 436). 

The following equations express the reaction for other nitrates of 
heavy metals. In all cases, the experimental observations are exactly 
as above except for the colours of the oxides, which are stated. With 
the nitrates given below there is no decrepitation. 

2Ca(N0 3 ) a 

. 2CaO + 4NO a + O a 



2Mg(N0 3 ) 2 -+ 2MgO + 4NO s + O, 
white magnesium 


2Zn(N0 3 ) 2 -»- 2ZnO + 4NO g + O, 
white zinc oxide 

(yellow when hot; 
white when cold) 

2Cu(NOa) a -*- 2CuO + 4N0 4 + O, 
green copper(II) oxide 

Mercury(II) and silver nitrates 

The oxides of mercury and silver are decomposed by heat, there- 
fore the nitrates of these two metals leave the free metals when heated. 

2AgN0 3 
Hg(NO a ) 2 

> 2Ag + 2NO g + O, 

* Hg + 2NO. 4- O t 

Questions on this chapter will be found on page 453. 

Chapter 34 

Oxides of Nitrogen 


(formerly called nitrous oxide) 

THIS compound is most conveniently prepared by the action of 
heat on ammonium nitrate (Fig. 144). 


Fio. 144. 
Preparation of dinitrogcn oxide. 

On heating, the ammonium nitrate melts and effervesces. Di- 
nitrogen oxide and steam are given off. Most of the steam is con- 
densed to water as the bubbles pass up the gas-jar. 
NH 4 N0 3 -»-N 2 + 2H 8 

Ammonium nitrate will explode, on heating, if the amount of 
material in the vessel becomes very small, but the reaction is quite 
safe unless this happens. If desired, the gas may be prepared by heat- 
ing any mixture of salts which, by double decomposition, will yield 
ammonium nitrate— e.g., a mixture of potassium nitrate and am- 
monium sulphate, finely ground. 




(NHJ^SO, + 2K.NO3 -*- 2NH 4 N0 3 + K,S0 4 
This mixture is less liable to explode. 

Test for dinilrogen oxide 

The gas rekindles a brightly glowing splint. 

It may be distinguished from oxygen by the following 

Dinitrogen oxide has a sweet and sickly smell (oxygen has 
no smell). 

Invert a gas-jar of dinitrogen oxide over cold water and 
shake it. The level of water in the gas-jar will rise, showing 
the gas to be fairly soluble in water. Oxygen is almost 
insoluble in water and no rise would be observed. 

Dinitrogen oxide does not give brown fumes with nitrogen 
monoxide. If a little oxygen is bubbled up into a volume of 
nitrogen monoxide enclosed over water, brown fumes of 
nitrogen dioxide are formed. 

Properties of dinitrogen oxide 

The gas is colourless, neutral to litmus and has a rather sweet, 
sickly smell. It can produce insensibility for short periods and is used 
as an anaesthetic for minor surgical operations, such as are required 
in dentistry. Insensibility lasts for a minute or two only. The period 
of insensibility can be prolonged if the gas is mixed with about 20% 
oxygen and inhaled, the oxygen being necessary to keep the patient 
alive. A trace of carbon dioxide must also be present to stimulate the 
respiratory centres and maintain breathing. Patients recovering from 
the effects of dinitrogen oxide may become hysterical; hence its com- 
mon name — 'laughing gas'. 

Density. Dinitrogen oxide has a density of 22 compared with 
hydrogen. (Molecular weight of dinitrogen oxide = 44.) 

Solubility. The gas is fairly soluble in cold water. 

Dinitrogen oxide as a supporter of combustion 

Action on a glowing splint. If the splint is glowing brightly, it will 
be rekindled by dinitrogen oxide, but, if only feebly glowing, it will 
be extinguished. 

To be rekindled, the glowing portion of the splint must be hot 
enough to decompose some dinitrogen oxide into nitrogen and 
oxygen. The mixture will then be rich enough in oxygen to stimulate 
the combustion of the splint and cause it to burst into flame. A 
feebly glowing splint will not be hot enough to decompose the 
dinitrogen oxide, and so will be extinguished, having no free oxygen 
with which to burn. 



candle wax 


Combustion of a candle, sulphur, carbon, and phosphorus 

These materials are placed on deflagrating spoons and made to 
burn by heating in a Bunsen flame. (The candle is, of course, merely 
lighted.) When plunged into dinitrogen oxide, they all burn more 
vigorously than in air, the candle with its familiar yellow flame, 
sulphur with a blue flame, and phosphorus with a bright yellow 
flame. The oxides of the burning elements are formed and gaseous 


" + 2N,0 -»■ CO, + 2N, 

+ N,0 -*■ H,0 + N, 
C + 2N,0 ->• CO, + 2N, 

S + 2N,0 -> SO, + 2N, 

4P + lONgO -> P«O 10 + ION, 
dense white 




Feebly burning sulphur may be extinguished by dinitrogen oxide. 
For the explanation of this, see the action of dinitrogen oxide on a 
glowing splint (above). 

Formula of dinitrogen oxide 
See p. 118. 


(formerly called nitric oxide) 

Nitrogen monoxide is produced, mixed always with other oxides 
of nitrogen, by the action of nitric acid on most metals. The common- 
est reaction used is that of moderately concentrated nitric acid on 
copper. Set up the apparatus shown in Fig. 145. Cover the copper 
with water and add concentrated nitric acid, about equal in volume 
to the water. Vigorous effervescence occurs and the flask is filled 
with brown fumes. These are nitrogen dioxide produced partly by 
the action of the acid upon the copper and partly by the oxidation of 
the main product, nitric oxide, by the oxygen of the air in the flask. 

2NO + O, -»- 2NO, 


The brown fumes dissolve in the water over which the nitrogen 



Cone, nitric acid and water 
{equal volumes) 





Fig. 145. 
Preparation of nitrogen monoxide. 

monoxide is collected as a colourless gas. A green solution of 
copper(II) nitrate is left in the flask. 

3Cu + 8HNO s -»■ 3Cu(N0 3 ) a + 4H a O + 2NO 
Tests for nitrogen monoxide 

Exposure to air. Remove the cover from a gas-jar of 
nitrogen monoxide. Reddish-brown fumes are at once pro- 
duced by oxidation of the gas by oxygen of the air. 

2NO + 0,-»-2N0 2 
Nitrogen monoxide is the only gas to give this action. 

Action on iron(II) sulphate solution. Prepare a cold solu- 
tion of iron(II) sulphate in dilute sulphuric acid. Pour it into 
a gas-jar of nitrogen monoxide. The dark brown or black 
coloration is caused by formation of a black compound, 

FeS0 4 + NO -> (FeS0 4 ).NO 

(The gas can be obtained in a pure state by heating this 

Properties of nitrogen monoxide 

Nitrogen monoxide is colourless and almost insoluble in water. Its 
density is 15 times that of hydrogen; it is slightly denser than air of 
which the density is 14.4. It is neutral to litmus. 

The smell of the gas is unknown because it combines with oxygen 
of the air (see tests above); it can never be collected by displacement 
of air. 

Nitrogen monoxide as a supporter of combustion 

Nitrogen monoxide will support the combustion of those burning 
materials whose flames are hot enough to decompose it and so 
liberate free oxygen with which the material may combine. A splint, 
candle, sulphur and glowing charcoal are all extinguished by the gas, 
but it supports the combustion of strongly burning phosphorus or 

2Mg + 2NO -»■ 2MgO + N, 

Formula of nitrogen monoxide. 

Seep. 117. 


Nitrogen dioxide is given off, together with oxygen, when nitrates 
of heavy metals are heated. The most suitable nitrate to use is leadfjl) 
nitrate, because it crystallises without water of crystallisation, which 
is found in crystals of most nitrates and which would interfere with 
the preparation. 

The lead(TT) nitrate decrepitates and melts on heating. It effervesces. 

lead (E) 



Fio. 146. 
Preparation of nitrogen dioxide. 

giving a brown gas, nitrogen dioxide, and oxygen. The nitrogen 
dioxide is liquefied in the freezing-mixture (Fig. 146) and collects in 
the U-tube as a green liquid (yellow if pure), oxygen passing on as gas 
and escaping. Lead(II) oxide remains in the tube as a yellow solid 
fused into the glass. 

2Pb(N0 3 ) 2 (c) -> 2PbO(c) + 4N0 2 (g) + O t (g) 



Properties of nitrogen dioxide 

Nitrogen dioxide is usually seen as a reddish-brown gas, though 
the boiling-point of liquid nitrogen dioxide (22°C) is above the usual 
atmospheric temperatures. It has a pungent, irritating smell and is a 
rather dangerous gas on account of its tendency to set up septic 
pneumonia if inhaled. It should not be allowed to escape in quantity 
into the open laboratory. 

Action of nitrogen dioxide with water 

Pour a little of the liquid nitrogen dioxide prepared above into a 
dry gas-jar in the fume cupboard and allow the liquid to vaporise. 
Add water and shake. The brown fumes dissolve leaving a pale blue 
liquid. Add to it blue litmus solution. The litmus is turned red. This 
is because the solution contains nitric and nitrous acids. 

2N0 2 -f- H 2 -> HNO s + HN0 8 
nitric nitrous 
acid acid 

The gas is a mixed anhydride. Similarly with sodium hydroxide, it 
gives a mixture of sodium nitrate and sodium nitrite. 

2NaOH + 2NO„ ->■ NaN0 3 + NaNO s + H 2 
sodium sodium 

nitrate nitrite 

Nitrogen dioxide as a supporter of combustion 

Like the other two oxides of nitrogen we have considered, nitrogen 
dioxide will support the combustion of a burning substance whose 
flame is hot enough to decompose it and supply free oxygen. It will 
support, for example, the combustion of strongly burning phosphorus 

2P 4 + 10NO, -► 2P 4 O 10 + 5N, 
but not that of sulphur or a candle. 

The use of nitrogen dioxide in the manufacture of sulphuric acid is 
discussed on p. 411. 

Dissociation of nitrogen dioxide 

At 150°C, nitrogen dioxide is very dark brown in colour. Its 
vapour density is 23, and its molecular weight is therefore 46, cor- 
responding to the formula NO, (N = 14; O = 16). As the temper- 
ature falls, the vapour density of the gas gradually increases, until at 
22°C it approaches 46, corresponding to a molecular weight of 92 
and a formula of N 8 4 . At the same time it becomes lighter in 

This must mean that, on heating, nitrogen dioxide dissociates, and, 
at any temperature between 22°C and 150°C, contains both N 4 0« 


and NOg molecules, the proportion of the latter increasing, as the 
temperature rises to 150°C, when only NO. molecules are present. 


N s O« ^ 2N0 2 

light yellow c°" 1 dark brown 

For a more complete consideration of dissociation, see p. 159. 


Fixation of atmospheric nitrogen 

The fertility of soil depends in part on the presence in the soil of 
certain chemical elements. These elements are potassium, nitrogen 
and phosphorus, together with traces of iron, sulphur and others. 
We shall consider, for the present, nitrogen alone. 

Every time a crop is taken from a given patch of soil, some of the 
nitrogen previously contained in the soil is removed in the form of 
complex organic compounds, which arc part of the tissue of the 
plant. This nitrogen was absorbed from the soil as dissolved nitrates 
by the roots of the plant and this is the only manner in which the vast 
majority of plants can absorb and use nitrogen. It is obvious that 
unless nitrogen is continually supplied to the soil to balance the loss 
suffered by removal of crops, the fertility of the soil will decrease and 
its yield become meagre. 

The soil receives some nitrogen by natural means. Certain plants, 
for example peas and beans, always have colonies of bacteria on their 
roots which are able to convert the nitrogen of the air into compounds 
which pass into the soil. Electrical discharges in the atmosphere, 
such as lightning, cause some slight combination of oxygen and 
nitrogen and this leads to the passage of nitrogen into the soil as 
nitrates, dissolved in rain water. This was a natural counterpart of 
Birkeland and Hyde's process, now obsolete. Nitrogen-fixing 
bacteria, living free in the soil, are another important agency supply- 
ing nitrogen to the soil from the air. Nitrogenous fertilisers are also 
used to make good the loss of nitrogen. These fertilisers fall into the 
following classes: 

1. Ammonium sulphate derived from the ammonia produced 
when coal is distilled. 

2. Sodium nitrate from the deposits of this substance in Chile— 
'Chile saltpetre', NaN0 3 . 

3. Farmyard manure. 

4. Fertilisers produced by manufacture, using the nitrogen of the 
air. These fertilisers include ammonium sulphate, for which the 


ammonia is prepared from nitrogen of the air by Haber's 
process (p. 431), and nitrates, also obtained from nitrogen of the 
air by oxidising ammonia to nitric acid. The nitrogen of the air 
is inexhaustible so that adequate supplies of nitrogenous fer- 
tilisers are now assured for ever. 

Though most plants can only use nitrogen as nitrates, it need not 
actually be supplied to the soil in this form. Bacteria in the soil will 
carry out the oxidation of any nitrogen compound to nitrates and the 
plants can then utilise it. 

The sum total of all these processes is called the 'Nitrogen cycle'. 
A simplified form of it is given in Fig. 147. 

Removed for food 
Loss of 


death and decay 


Absorbed Bacteria 

by or 

roots lightning 

in the soil "* 


NaNQ 3 

in the soil 

Atmospheric nitrogen 

Haber's process 

and oxidation 

ofNH 3 

Nitrates from nitrogen 
of the air 

Ammonium sulphate 
from coal 
Farmyard manure 

-►■Ammonium compounds 
in the soil 




oxidising bacteria 

in the soil 

Fig. 147. 
The nitrogen cycle. 

Haber's process for fixation of atmospheric nitrogen 

Fixation of atmospheric nitrogen by Haber's process has already 
been considered (p. 431). 

The product, ammonia, may be converted to ammonium sulphate 
(p. 417) or oxidised by passing the ammonia with excess of air over 
platinum wire gauze (catalyst) at red heat 

4NH 3 + 50, -»• 4N0 + 6H a O 



The nitrogen monoxide so formed is rapidly cooled and combines 
with the oxygen from excess of air to form nitrogen dioxide. 

2NO + O, -*• 2NO, 
The nitrogen dioxide, in the presence of more air, is then absorbed in 
hot water, the conditions being chosen to yield nitric acid by the 

2H t O + 4NO„ -r- O s -> 4HNO s 
A valuable fertiliser can be obtained by neutralising the acid with 
lime and the calcium nitrate, mixed with excess of lime to form a non- 
deliquescent basic salt, is applied to the soil. 

Ca(OH) 2 + 2HNO s ->- Ca(N0 3 ), i 2H t 


1 Describe the usual laboratory method of making nitric acid, and give 
an account of its action on (a) one base; (ft) one metal; (c) one non-metallic 
element; id) one carbonate. (B.) 

2 Describe shortly how you would obtain, starting with nitric acid, 
specimens of (a) oxygen; (b) nitrogen monoxide; (c) dinitrogen oxide; and 
(d) nitrogen dioxide. (N.U.J.B.) 

3. Give equations illustrating the effect of heat upon nitrates. 
Describe concisely the properties of the gaseous products of these 

reactions. (O. and C.) 

4. Describe the preparation and collection of nitrogen monoxide in the 
laboratory and give an account of its chief properties. How would you 
prove that it consists of nitrogen and oxygen? 

Explain, very briefly, how it is formed in the earth s atmosphere, and 
indicate how the acid subsequently produced becomes 'fixed' in the soil as 
nitrate. (B.) 

5 How would you prepare dinitrogen oxide? Give an account of the 
properties and uses of this gas. Draw and describe the apparatus you 
would use to show that dinitrogen oxide contains its own volume of 
nitrogen. (D.) 

6 Describe the preparation and properties of nitrogen monoxide. Three 
litres of a mixture of nitrogen monoxide and dinitrogen oxide were passed 
over red-hot copper, and 2.2 litres of nitrogen collected. Calculate the 
composition of the mixture. (O. and C.) 

7. How would you prepare and collect some nitrogen monoxide free 

from nitrogen? . 

Describe and account for what you would observe during the pre- 
paration. What happens when this gas is (a) passed into ferrous sulphate 
solution, and (b) mixed with oxygen? (N.U.J.B.) 

8. Given ammonium nitrate how would you proceed to obtain a sample 
of (a) dinitrogen oxide; (A) ammonia? . 

How would you obtain from ammonia a sample of nitrogen .' wnat is 
the reaction between an aqueous solution of ammonia and (a) iron(III) 
chloride solution; (b) dilute sulphuric acid? (C.W.B.) 



9. How is nitrogen usually prepared and collected in the laboratory? 
How is nitrogen obtained industrially? 

Outline the industrial method by which nitrogen may be converted into 
(a) ammonia; (6) nitric acid. 

How and under what conditions does copper(II) oxide react with 
ammonia? (N.U.J.B.) 

10. You are provided with concentrated sulphuric acid, slaked lime, 
ammonium chloride, potassium nitrate, water, a source of heat and any 
apparatus needed, but no other chemicals. Explain briefly how you would 
prepare (a) nitric acid, (b) ammonia solution, (c) quicklime, {d) pure dry 
calcium sulphate. (N.U.J.B.) 

11. Describe one experiment in each case to show that ammonia gas is 
(a) very soluble in water and gives an alkaline liquid, (b) a reducing agent. 
How, and in what conditions, does ammonia react with (/) chlorine, (//) 
hydrogen chloride? How would you separate a sample of dry ammonia 
gas from a mixture containing this gas and nitrogen ? 

12. Outline the Haber process for the fixation of atmospheric nitrogen. 
How is the product of this fixation converted to two materials which can be 
utilised as fertilisers for agricultural land? 

13. State the volume composition of ammonia gas. Assuming the 
diatomicity of nitrogen and hydrogen, deduce the molecular formula of 
ammonia, stating any principle of which you make use. How may am- 
monia be converted into (a) nitrogen; (b) nitrogen monoxide; (c) dry, 
crystalline ammonium chloride? 

Chapter 35 



PHOSPHORUS is a very active non-metal and is not found un- 
combined in nature. It was first isolated by Brandt, a Hamburg 
chemist, in 1669, who obtained it by distilling concentrated urine 
with sand. It is extracted nowadays from bones and mineral phos- 
phates. The residue left after burning away the organic matter from 
bones contains calcium phosphate, Ca 3 (PO,) s . This substance also 
occurs as the mineral, apatite. 


A charge of calcium phosphate, sand and coke is fed contin- 
uously into an electric furnace, Fig. 148. At the very high temper- 
ature obtained (the electric current is merely to produce the high 

carbon _ 


carbon electrode 

Fig. 148. 

Extraction of phosphorus. 



temperature; there is no electrolysis) the following reaction takes 

2Ca 3 (P0 4 ) a + 6SiO a + 10C-* 6CaSiO a + 10CO + P 4 

calcium silicate 

The phosphorus distils over with the carbon monoxide and the 
vapour is led below the surface of water, when the phosphorus 
solidifies to a white solid. The calcium silicate is formed as a molten 
slag which is run off from time to time at A. 

Chemistry of the action. Most salts can be considered as being made up 
from a basic oxide and an acidic oxide. 

2Ca,(P0 4 ), = 6CaO.P 4 O I0 
Silica is also an acidic oxide, hence 

6CaO + 6SiO, — >- 6CaSiO, 
calcium silicate 
(cf. calcium carbonate) 

The carbon (a reducing agent) reduces the oxide of phosphorus to the 

P 4 O 10 + 10C-»-P 4 + 10CO 
On adding these equations together the main equation will be obtained. 


Phosphorus exists in two chief allotropic forms, white or yellow 
phosphorus and red phosphorus. The latter is the stable form. (See 
p. 457.) 

Yellow phosphorus is a white solid which becomes pale yellow on 
exposure to light It is of density 1.8 g/cm 3 and is soluble in carbon 
disulphide. It smoulders in air owing to oxidation and this action 
causes it to glow in the dark. 

P 4 + 30 8 -»-P 4 8 
It is usually kept below the surface of water. Great care should be 
taken in handling it, for it gives off a very poisonous vapour and catches 
fire very readily. It burns in air or oxygen, giving off white fumes of 
oxides of phosphorus. This variety of phosphorus is formed when 
the vapour of phosphorus is suddenly cooled. 

Red phosphorus is the stable variety at all temperatures. It is not 
poisonous as is yellow phosphorus, and does not catch fire so readily. 
It can be made by heating yellow phosphorus in an inert atmosphere 
(usually with a little iodine to act as catalyst) to a temperature of 
about 250°C. It is insoluble in carbon disulphide. 

Experimental evidence that the two allotropes are chemically 
identical is the following. 1 g of white phosphorus heated to about 
250°-300°C in an inert atmosphere yields exactly 1 g of red phos- 



phorus and nothing else. Further, if equal weights of red and white 
phosphorus are converted fully to any given phosphorus compound, 
exactly the same weights of product are obtained. 

Comparison of allotropes of phosphorus 

Red White (or yellow) 

Opaque red solid 

Density 2.3 g/cm 3 . 


Sublimes at 400°C. 

Insoluble in carbon disulphide. 

Ignites in air at 260°C. 

No action with hot sodium 
hydroxide solution. 

Unoxidised at ordinary tempera- 
tures in air. 

Colourless translucent solid 

(turns yellow). 
Density 1.8 g/cm 3 . 
Very poisonous. 
Melting-point 44°C. 
Soluble in carbon disulphide. 
Ignites in moist air at 30°C. 
Forms phosphine with hot 

sodium hydroxide solution. 
Rapidly oxidised at ordinary 

temperatures in air. 

White phosphorus is the unstable variety. It has the higher vapour 
pressure and is very slowly reverting to red phosphorus in room 

Phosphorus is used in the match industry, in rat poisons and in 
making smoke bombs. About 80% of it is converted to phosphoric 
acid, H 3 P0 4 . 

Match industry 

In earlier days, matches contained white phosphorus and an 
oxidising agent and the mixture was caused to burn by rubbing on 
sandpaper. The use of white phosphorus in industry caused hundreds 
of work-people to suffer from phosphorus poisoning, which took the 
form of the rotting of the bones of the face and jaw ('phossy-jaw'). 
Its use was then forbidden by law and matches nowadays consist of 
compounds of phosphorus (sulphides as a rule) and oxidising agents. 
The friction of rubbing on a rough surface generates enough heat to 
start the combustion. In 'safety' matches the phosphorus (red variety) 
is on the side of the box and thus the match-head, which contains the 
oxidising agent, is useless without the box. 

Chemical properties of yellow phosphorus 

Phosphorus will combine readily with oxygen, chlorine and sul- 
phur. It is a reducing agent and readily attacks oxidising agents with 
the formation of oxides of phosphorus which are soluble in water, 



forming acids. It is attacked by alkalis forming phosphine. Phos- 
phorus forms two series of compounds, exhibiting valencies of 3 and 
5. For phosphorus in periodicity, see pp. 84-90. 

Preparation of phosphine (phosphoretted hydrogen), PH 3 

Fit up the apparatus as shown in Fig. 149 in a fume-chamber. 
Place sodium hydroxide (caustic soda) solution and a few small 
pieces of yellow phosphorus in the flask, sweep out the air by means 
of coal-gas and warm the mixture. A gas is given ofi", phosphine, and 
on coming into contact with the air it ignites spontaneously (the 

sodium \ 
hydroxide \ 
solution (20%) 
and yellow 


Fio. 149. 
Preparation of phosphine. 

ignition is due to an impurity) and forms white vortex rings of oxides 
of phosphorus. 

3NaOH + P 4 + 3H a O-^ SNaHjPOj + PH 3 

sodium phosphine 

Phosphine is a colourless gas with a garlic-like odour and is ex- 
tremely poisonous. It shows similarities to ammonia, forming salts, 
for example, PH 4 Cl,phosphonium chloride (cf. ammonium chloride). 

Phosphorus trichloride, PC1 3 

This is made by allowing dry chlorine to react with warm yellow 
phosphorus in an inert atmosphere. 

P 4 + 6CI 2 -*■ 4PC1 3 

Phosphorus trichloride is a liquid (boiling-point, 76°C) which 



fumes in air owing to the action of moisture upon it. The liquid is 
attacked by water, forming phosphorous acid. 

PC1 3 + 3H,0 -» H 3 PO a + 3HC1 

Phosphorus pentachloride, PC1 6 

This is made from phosphorus trichloride by the action of chlorine 
upon it. A yellowish solid separates out. It sublimes when heated, 
decomposing into the trichloride and chlorine, but on cooling forms 
the original pentachloride. 




PC1 3 + Cl» 

phosphorus phosphorus chlorine 
pentachloride trichloride 
It also attacks water vigorously forming phosphoric acid, 
Pd 5 + 4H a O -*■ H 3 P0 4 + 5HC1 

Phosphorus trioxide (phosphorous oxide), P 4 0„ 

This is made by burning phosphorus in a limited supply of air and 
passing the mixture of trioxide and pentoxide through a tube sur- 
rounded by a water jacket at 50°C, and containing a loose cotton- 
wool plug. The pentoxide remains solid and is retained by the plug. 
The trioxide vapour passes on and is condensed in a freezing- 

P« + 30 2 -»-P«0 4 

P« + 5O 1 -»-P«O l0 
Phosphorous oxide is a white volatile solid which readily reacts 
with water to form phosphorous acid. 

P 4 0, + 6H a O ->- 4H 3 P0 3 

Phosphorus pentoxide (phosphoric oxide), P 4 O 10 

This is made by igniting phosphorus in a plentiful supply of air A 
small piece of phosphorus is placed in a crucible on a plate of glass 
under a bell-jar containing dry air. The phosphorus is ignited. It 
burns with a brilliant flame and a white solid finally settles on the 
plate. This is quickly scraped into a dry bottle. 
P 4 -t-5O s -*P 4 O 10 



It can be purified by heating it in a current of dry oxygen. The 
latter converts any trioxide formed into pentoxide. 

Phosphorus pentoxide is a white solid which reacts vigorously with 
water, and is one of the best drying agents known. With hot water it 
forms phosphoric acid. 

P«O I0 + 6H,0 -*■ 4H3PO4 

Acids of phosphorus. Many acids of phosphorus exist, the chief of 
which are phosphorous and phosphoric acids. 

Preparation of phosphoric acid, H 3 P0 4 , from phosphorus 

Heat a small amount of red phosphorus in a dish with fairly con- 
centrated nitric acid (1:1, acid and water) for some time on a water 
bath in a fume-chamber. Dilute with water and filter off any un- 
attacked red phosphorus. The product is a solution of phosphoric 
acid, H3PO4. It may be concentrated, and any nitric acid removed, 
by heating it till the temperature reaches 200°C, above which it 
decomposes. The nitric acid is reduced to oxides of nitrogen and 
there is a copious evolution of brown fumes during the experiment. 
The presence of a phosphate can be shown by warming a portion of 
the filtrate with nitric acid and excess ammonium molybdate. A 
yellow coloration or precipitate of ammonium phosphomolybdate 
indicates the presence of a phosphate. 

P 4 + 2OHNO3 -> 4H,PO« + 20NO a + 4H a O 
phosphoric nitrogen 
acid dioxide 

This acid, H 3 P0 4 , should strictly be called orthophosphoric acid; its 
dehydrated product, HP0 3 , is metaphosphoric acid. 

Phosphatic fertilisers 

Phosphorus is an element essential to soil fertility. Plants absorb it 
from the soil and, if the crop is consumed by man, phosphorus 
passes into his bone structure and protoplasm. Much of it is lost to 
the soil in general by sewage disposal and by depositing human 
remains in cemeteries. 

The loss must be made good. Bone-meal and basic slag are used, 
both of which contain calcium phosphate, Ca^POJj. This com- 
pound is, however, almost insoluble in water and becomes available 
to the plants very slowly. A more soluble material, quicker in action, 
can be made by stirring calcium phosphate with an appropriate 
weight of 65% sulphuric acid. An acid calcium phosphate is formed. 

Ca 3 (P0 4 ) 2 + 2H a S0 4 ->- Ca(H a P0 4 ) a + 2CaSO, 
The product is dried and sold as 'superphosphate'. 



Electrothermal reactions 

The extraction of phosphorus (p. 455) is an electrothermal reaction, 
that is, the current is used to produce heat and not for any electrolytic 
purpose. In such reactions, alternating current may be used; for 
electrolytic decompositions, direct current is essential. 

Calcium carbide is manufactured electrothcrmally. A steel furnace 
is used, lined with carbon, which acts as one of the electrodes. The 
other is a carbon rod, placed vertically. The furnace contains coke 
(carbon) and quicklime (calcium oxide). Alternating current causes 
arcing between the pieces of coke; this raises the temperature to 
about 2000°C. 

CaO + 3C -»■ CaC a + 2CO 
Calcium carbide is produced molten, later cools to form a solid and 
is broken up and graded in an atmosphere of nitrogen to prevent 
air-acetylene explosions. 

Calcium carbide is important chiefly because it yields acetylene, 
C a H a , with cold water. 

CaC a + 2H a O -► Ca(OH) a + C a H a 
Acetylene is important as a source of polyvinyl chloride (PVC) 
plastic. Hydrogen chloride reacts additively with acetylene to give 
vinyl chloride. 

HC3=CH + HC1 -> H,C=CHC1 
(the group H a C=CH— is called vinyl) 

Vinyl chloride can be polymerised (in aqueous emulsion and initiated 
by peroxides) to PVC, which consists of long molecular chains of the 


. . . _CH s -CHCl-CH a -CHCl-CH a -CHCl-CH a -CHCl- 
. . ._CH a -CHCl-CH a -CHCl-CH a -CHCl-CH a -CHCl-. . . 

(Compare the polymerisation of ethylene to polythene, p. 329.) PVC 
is a thermoplastic much used in electrical insulation and water- 
proofing. Notice that the double bonds (unsaturation) are lost in 
polymerisation and PVC is a saturated compound. 


1 Describe a method (a) for the preparation of white phosphorus from 
calcium phosphate; (b) for the conversion of white phosphorus into the red 

° Compare and contrast the properties of these two forms of phosphorus. 
What reaction, if any, has each form with chlorine, and with sodium 
hydroxide solution? (D.) 



2. Compare the properties of the two forms of phosphorus. How may 
phosphorus be converted into two of the following: (a) phosphorus penta- 
chloridc; (6) phosphine; (c) phosphorus pentoxide? (O. and C.) 

3. Describe with necessary detail how you would prepare from phos- 
phorus reasonably pure specimens of (a) orthophosphoric acid ; (b) phos- 
phorus trichloride. (O.) 

4. Show by giving three chemical properties in each case that sodium 
is a metal and that phosphorus is a non-metal. 

Starting with metallic sodium and yellow phosphorus, describe fully how 
you would prepare specimens of (a) phosphine, and (b) sodium phosphate. 

5. How does phosphorus react with (a) oxygen; (b) chlorine; (c) sodium 

Describe how the products behave when brought into contact with 
water. (O. and C.) 

6. Describe the preparation of (a) phosphorus trichloride, and (b) 
phosphine. Show how the latter is related to ammonia. (L.) 

7. How is phosphorus obtained from bone ash? State how white 
phosphorus can be converted into red phosphorus, and name three 
respects in which they differ from one another. Write the formula; of two 
chlorides, one of hydride, and two oxides of phosphorus. How does the 
highest oxide of phosphorus react with water and for what purposes is this 
oxide used ? (L.) 

8. Contrast, tabulating your answer, six properties of two allotropes of 
phosphorus. Starting with one of these forms, briefly describe the pre- 
paration of the other form from it. How is phosphine usually prepared in 
the laboratory? Give details. (L.) 

Chapter 36 

Metals; Extraction and Uses; 
Non-Metals Compared 

UNTIL recently, physical differences between metals and non- 
metals have been regarded as quite significant. For example, 
gold and silver, with high density, lustre, malleability and ductility, 
have been regarded as two of the most typical metals; sulphur, which 
has a much lower density and no lustre, and is very brittle, was 
regarded as a typical non-metal. This outlook has been largely 
replaced by the following: 


A metal is now defined as an element which can ionise by electron 
loss. The number of electrons lost per atom is the valency of the metal 
and the ion carries an equal number of positive charges, as: 
Na — e - — >■ Na + (univalent) 
Mg-2e~-»-Mg 2+ (divalent) 
Al — 3e~— *■ Al 3 + (trivalcnt) 
The following are important chemical properties of metals which 
are derivative from this ionisation behaviour. 

1. The oxide of a metal is basic, and, if soluble in water, gives an 
alkaline solution. 

This property follows from the fact that a typical metallic oxide 
is formed by oxygen accepting electrons from the metal as it ionises, 
to form the ion, O 2- . For example, 

Ca-»-Ca a+ + 2ei 
iO g + 2e--*-0 2 - 

or adding these, Ca + iO, -*■ Ca» + 0*~ 

If the oxide dissolves in water, it gives the reaction: 

O 2 - + H,0 e* 20H- 
so providing the hydroxyl ion, which is characteristic of alkalinity. If 





insoluble in water, the oxide does not produce alkalinity but acts as 
a basic oxide. 

This property arises from the behaviour of the O 8- ion or OH~ ion 
with hydrogen ion, H + , which is characteristic of acidity. The 
behaviour is: 

O 2 - + H,0 f* 20H-"! 
20H- + 2H + ^2H 8 OJ 

The second of these reactions is the essential nature of neutralisation, 
i.e., the formation of water from its ions, H + and OH - . The ions, 
Ca 2+ and 2CI~, derived from calcium oxide and hydrochloric acid 
may remain as the salt, calcium chloride, and correspondingly for 
other oxides and acids. The complete reaction can be represented as: 

Ca* + 8 - + H a O ^ Ca a+ (OH-) 2 
Ca a +(OH-), + 2(H + C1") -> Ca 2+ (Cl-), + 2H 2 

2. A metal can replace H+ ion in an acid and so produce a salt 

This replacement may occur directly or indirectly. If directly, the 
metal ionises, liberating electrons; the electrons are accepted by H + 
ions in the acid. These ions become hydrogen atoms, pair off as 
molecules and are liberated as gas. The metallic ions pass into 
solution and, in conjunction with the negative ions of the acid, CI~, 
S0 4 2_ etc., constitute a salt. For example : 

Zn-».Zn 3 + + 2e-\ 

2H + + 2e~ -> H 2 J 

For dilute sulphuric acid, 

Zn + (H + ),S0 4 *- ->- Zn s+ SO« 2 - + H, 

Indirect replacement occurs chiefly through the medium of the 
metallic oxide or hydroxide, as: 

Cu + K) t -»-Cu* + 2 - 
O*- + H,0 5* 20H- 
20H- + 2H+ ^ 2H,0 
If the acid is dilute sulphuric, Cu 2+ ions are left in solution, with 
S0 4 2 ~ ions, as the salt, copper sulphate, the other product being water. 

3. Metals form elect rovalent chlorides. 

Metallic chlorides are typically electrovalent in type because, in 
their formation, a metal loses electrons, which are accepted by 
chlorine atoms, e.g., 

Ca ->- Ca 2+ + 2e-\ 
CI, + 2e- ^ 2CI- / 

or adding these, Ca + CI, — >- Ca 2+ (C1-), 



Such chlorides are electrovalent, electrolytes when molten or in 
aqueous solution, and solids of very low volatility. 

4. Metals form few compounds with hydrogen 

This situation arises because hydrogen forms compounds most 
readily by covalency (i.e., electron sharing) or, in acids by electron 
loss, to form the ion, H + . Neither of these modes of combination is 
well suited to metals, which tend to operate by electron loss. 

A few very powerful metals (Na, K, Ca) can force hydrogen to 
accept electrons. 

The hydrides so formed are salt-like solids (compare the chlorides 
above) and liberate hydrogen with cold water, e.g., 
Na + ±H,-»-Na + H- 
H- + H 2 0-*OH- + H, 

When molten, they act as electrolytes and liberate hydrogen at the 

H--e--*iH a 
(to anode) 

In aqueous solutions, containing the usual ion, H + , hydrogen appears 

at the cathode during electrolysis. 

H+ + e-->iH, 

5. Metals are reducing agents 

This follows from the modern definition of a reducing agent (p. 
162) as an electron donor, e.g., 

K-»-K+ + e- 
Zn ->- Zn 2+ + 2e~ 

From the theoretical angle, the most characteristic physical pro- 
perty of a metal is its ability to conduct electricity, i.e., to pass a 
stream of electrons through itself if a potential difference is applied 
across it. This arises because metallic atoms part with electrons easily 
and a mass of metal always contains comparatively loose electrons. 
These move readily through the metal and are steadily replaced from 
the source of potential difference to maintain a flow of current. Con- 
duction of heat, lustre, malleability, ductility and high tensile 
strength are other physical properties possessed by some metals. They 
are very useful in practice but are not now accorded their former 
importance. Many metals have high density, but the most charac- 
teristic common metals, i.e., those which ionise most readily by 
electron loss, are sodium and potassium, and they have densities 



below that of water. They are also soft enough to be cut with a pen- 


If a non-metal forms ions, it does so by electron gain, i.e., the element 
is electronegative. The number of electrons gained per atom is the 
valency of the element for this purpose. (It may also exercise other 
kinds of valency.) The ion formed carries the corresponding number 
of negative charges but they never exceed two. 

|CI 3 + e~ -*■ CI- (univalent) 

£Br 8 + e~ —*■ Br - (univalent) 

±b 8 -f 2e~ — >- O 8 - (divalent) 

S + 2e--»-S a - (divalent) 

The following arc important chemical properties of non-metals 
which are connected with their tendency towards electron gain, 

1. The oxide of a non-metal is never basic. Its characteristic oxide is 
acidic. Its other oxides are acidic or neutral 

As explained on p. 464, a basic oxide is formed when an oxygen 
atom accepts electrons from a metallic atom to give the ion, O 8- . 
Since a non-metal is an acceptor of electrons, it cannot form an oxide 
in this way. The characteristic oxide of a non-metal, i.e., the oxide in 
which the non-metal exercises its maximum valency, is a covalent 
oxide, produced by formation of shared electron pairs. 

When combined with water, it forms an acid, e.g., 








Characteristic oxide: its action with water 

co, + h 8 o ^ h 8 co, ^ 2H+ + co,»- 

N a O B + H 8 ?s 2HN0 3 ^ 2H + + 2NO a - 

S0 3 + HgO ?± H 3 SO« ^ 2H+ + S0 4 2 - 

C1 2 7 + H a O ^ 2HCI0 4 ^ 2H+ + 2C10«- 

The H ion, characteristic of acidity, is produced in all these cases. 
When the non-metal exercises a reduced valency, the oxide is 
sometimes acidic, sometimes neutral, but never basic. Examples: 





N 8 
N a O« 
S0 8 
C1 2 

Nature of oxide 


2. A non-metal never replaces hydrogen in an acid to form a salt (as a 
metal does) 

Replacement of hydrogen in an acid arises from the acceptance by 
H + of electrons supplied by a metallic atom: 

2H + + 2e~ -*■ H 2 
A non-metal is an electron acceptor and so cannot bring about the 
above replacement. 

3. Non-metals form covalent chlorides 

The behaviour of the non-metal, phosphorus, in forming its 
trichloride is typical. 



CI* P ° 

U1 o r o 



CI— P 


• electrons from CI 
O electrons from P 

Each CI atom has also seven more electrons (not shown) in its outer 
electron layer. A covalent chloride of this kind is usually a volatile 
liquid, a non-electrolyte, and rapidly hydrolysed by water, as: 

PCl a + 3H g O -> H 3 PO s + 3HC1 
These properties are characteristic of non-metallic chlorides (though 
CC1 4 is not hydrolysed by water). Non-metals cannot supply electrons 
to form electrovalent chlorides. This is the behaviour of metals 
(p. 464). 

4. Non-metals combine with hydrogen to form many covalent hydrides 
These hydrides are formed by electron sharing, as: 
(Each bond represents one shared pair of electrons) 

H— C— H 








hydrogen sul- 
phide (covalent 

h— a 




hydrogen chlorid 
(covalent form) 

Simple hydrides of this type have the properties of covalent com- 
pounds; they are gaseous, form no ions when anhydrous and so are 
non-electrolytes. Salt-like hydrides are formed by metals (p. 465) but 
non-metals cannot supply the electrons required for this. 



5. Non-metals are oxidising agents 

This follows from the modern definition (p. 162) of an oxidising 
agent as an acceptor of electrons, e.g., 

iCl s + e--*-C\- 
S + 2e~ -»- S a ~ 

The most important physical property of non-metals is their very 
low electrical conductivity. Because of the tendency of a non-metal to 
attract electrons, a mass of it contains an absolute minimum of free 
electrons. This being so, no considerable electron flow can occur 
through the non-metal and it tends to be a very poor electrical 
conductor. (Graphite is exceptional, conducting electricity quite 
well.) In addition, non-metals are usually brittle and do not give sheets 
or wire. They have low tensile strength and no lustre, and cannot be 

To summarise in tabular form : 



Metallic oxides are basic and 
form alkalis if soluble in water. 

Metals replace hydrogen in acids 
forming salts. 

Metallic chlorides are electro- 
valent salts, electrolytes con- 
taining Cl~ ions. 

Metals form few stable hydrides. 
Na, K and Ca form salt-like 
hydrides, which contain ions, 



Characteristic oxides of non- 
metals are acidic; other oxides 
are acidic or neutral. 

Non-metals do not form salts in 
this way. 

Non-metallic chlorides are cc- 
valent, non-electrolytes, and 
are generally hydrolysed by 

Non-metals form many stable 
hydrides; they are covalent, 
the simple ones being gases 
(like NH 3 or CHJ and non- 
electrolytes if water-free. 

Metals are : 

Physical Properties 

Non-metals are : 

good conductors of electricity 
and heat; malleable (can be 
beaten into sheets); ductile 
(can be drawn into wire); 
lustrous and can be polished. 
Some metals are very dense and 
have high tensile strength. 

generally bad conductors of 
electricity and heat (graphite 
conducts electricity well); gen- 
erally brittle, not malleable or 
generally not lustrous. 



wrt . 



K tk n la Hi M CataAuCu T» 

Fio. 150. 

Melting-points of the common metals (excluding mercury). The metals are 
represented by their symbols. 

The following table shows an important relation between certain 
properties of metals and non-metals. 



Metals form oxides 

which are 

A few basic oxides 

are soluble in water 

and form 



Non-metals form 

oxides which are 


Acidic oxides 

combine with 

water forming 


basic oxides > + acids 






Extraction and the electrochemical series 












Very reactive. Never found as free element. Extracted by 
'electrolysis. All isolated after 1800. 

Not very reactive. May be found in nature as the free 
element. Known for a very long time. 

In this chapter the physical properties of the metals are arranged 
in the same order throughout. The electrochemical series shown 
above indicates (in a very rough way only) the inverse order in which 
the elements were isolated. Thus metals low in the series such as 
gold, silver and lead have been known since very early times. Metals 
high in the series proved very difficult to isolate and it was Davy's 
work on electrolysis which led to the isolation of potassium, sodium, 
calcium, magnesium and aluminium over a period of years from 1807 
(when Davy isolated potassium and sodium) to about 1850 (when 
aluminium was isolated). 

Aluminium was not obtained in the first place by electrolysis 
although its manufacture nowadays is entirely confined to that 
method, but it was isolated by the action of the very active element 
potassium on aluminium chloride. 

In view of what we have learnt about the electrochemical series, 
this order seems quite natural for we should expect the compounds 
of the very active elements such as sodium and potassium to be very 
stable substances. 

Metals low down in the scries are frequently found as the free 
elements, although they may also be obtained from ores because the 
amounts found as the free metal are not sufficient for industrial 
purposes. Gold, however, the last clement of the series, is found and 
mined almost entirely as the free element. 

Extraction is a reduction 

Metals are found as a rule as their oxides or salts, chiefly sulphides 
and carbonates, which are electrovalent compounds and contain the 



metal in ionised condition. During extraction, the metallic ion takes 
up the necessary number of electrons to convert it to the correspond- 
ing atom. This process of electron gain is a reduction and the electrons 
are supplied by the reducing agent concerned in the reaction. For 
example, in the ore, galena, Pb 2+ .S*~, lead exists (as shown) in ionic 
form. After extraction, it exists as atoms of lead. Therefore, at some 
stage, the change: 

Pb 2+ + 2e~ — *■ Pb (a reduction) 

must have taken place. The electrons arc supplied when the reducing 
agent, carbon, acts upon lead oxide, produced as an intermediate 
product in the extraction. 

O 2 - + C -»- CO + 2e- 

The reduction in the case of the first members of the series is 
brought about by electrolysis, which can be regarded as a very power- 
ful oxidation and reduction process which often results in the forma- 
tion of the elements present in the compound electrolysed (see 
Electrolysis, p. 133). As far as the metals are concerned the process 
is one of reduction, although no reducing agent (in the ordinary sense 
of the term) is used. 

The cathode acts as a reducing region by supplying electrons. The 
usual reducing agents employed for the less electropositive metals are 
carbon in the form of coke, and carbon monoxide (made from the 
coke by passing a limited amount of air over the hot coke). You may 
ask why the usual laboratory reducing agents are not employed ; for 
example, hydrogen. The reason is that, in an industrial process the 
chief concern is cost, and coke easily wins on that score. 



These metals occur chiefly as the chloride, sodium as common salt 
in the huge salt deposits of Cheshire, and elsewhere, and potassium 
chloride as carnallite (together with magnesium chloride) in the 
deposits in Stassfurt (see p. 376). Other sources of the elements are: 

sodium carbonate 
sodium nitrate 
sodium chloride 
potassium nitrate 
potassium carbonate 

in Africa and in ash of sea plants. 

Chile saltpetre in Chile. 

in sea-water and many salt lakes. 

saltpetre found in India. 

in the ash of land plants. 


Potassium is obtained by the electrolysis of its molten hydroxide. 



Physical properties of potassium 


Tensile strength 

Conduction of heat 
and electricity 


White metal, possessing a lustre. 

0.86 g/cm- 3 . 

Malleable and ductile. 

Does not possess tensile strength to any 

appreciable extent 

Good conductor of heat and electricity. 

In chemical properties potassium is very similar to sodium, but it is 
slightly more reactive. Thus, when a small piece of potassium is 
placed on water it darts about, melts and gives off hydrogen which at 
once ignites and burns with the oxygen of the air, giving a lilac- 
coloured flame. 

2K + 2H 2 -»• 2KOH + H, 

Flame tests for potassium (see p. 374) 

Compounds of potassium (especially the chloride) colour the 
Bunsen flame lilac and the colour is still visible through blue glass. 

Hence, although the lilac colour of potassium is easily 
masked by the presence even of traces of sodium, the 
potassium colour is visible when viewed through blue glass. 

In general, the salts of potassium are less soluble in water than the 
corresponding salts of sodium (the principal exceptions are potassium 
hydroxide, potassium carbonate and hydrogen carbonate). This 
accounts for the use in medicine of such substances as potassium 
permanganate, potassium iodide, potassium chlorate, where one 
would expect the cheaper sodium salt to be used. The extra cost of 
preparing the potassium salt instead of the cheaper sodium salt is 
outweighed by the ease of obtaining the potassium salt from the 
solution because of its lower solubility. 

Extraction of sodium 

Sodium was first isolated in 1807 by Davy. He electrolysed caustic 
soda between platinum electrodes, obtaining small beads of sodium 
metal at the cathode. A more suitable version of this extraction is 
illustrated by Fig. 151. Note that the electrolyte is molten caustic 
soda, i.e., solid caustic soda heated till it liquefies. If water were 
present, sodium could not be produced. The caustic soda (sodium 
hydroxide) gives the ions, Na* and OH". 

At the cathode 
Na+ + e - _>. Na 


At the anode 
OH- - e~ -»- (OH) 
Then: 4(OH) -»■ 2H t O + O, 


Caustic Soda 

Causae Soda 

Fig. 151. 
Preparation of sodium. 

Industrial extraction of sodium (Downs process) 

In this process, common salt is electrolysed in the molten condition. 
As the melting-point of the salt is high (about 800°C), calcium 

Sodium chloride odded 

Fused sodium 
chloride and 
calcium chloride 



Iron gauze 


FlQ. 152. 

Extraction of sodium from common salt. 



chloride is added to lower the melting-point; it becomes about 
600°C. The Downs cell (Fig. 152) has an outer iron shell, lined with 
firebrick. A diaphragm of iron gauze screens the carbon anode from 
the ring-shaped iron cathode that surrounds it. Chlorine escapes via 
the hood. Sodium collects in the inverted trough, T, placed over the 
cathode, rises up the pipe, P, and is tapped off through the iron 
vessel, V. Sodium chloride produces ions Na + and Cl~. 

At the cathode At the anode 

Na + + e--*Na C\--e~— »-CI 

(a reduction) (an oxidation) 

Then: CI + CI -+ CI, 

Physical properties of sodium 




White, silvery, shining metal. Rapidly 



0.97 g/cm- 3 . 


Very malleable. Can be cut with a knife. 

Tensile strength 

Does not possess tensile strength to any 

appreciable extent. 



Conduction of heat 

and electricity 

Conducts both heat and electricity. 

Action of sodium on exposure to air 

Sodium is attacked by the oxygen of the air to form sodium oxide. 
The moisture present combines with some of the oxide to form the 
hydroxide and finally, after a time, the carbon dioxide of the air com- 
bines with the sodium hydroxide to form sodium carbonate. 

4Na + 0„-*-2Na 1 
Na 2 + H a O -»• 2NaOH 
2NaOH + C0 2 ->- Na 2 C0 3 + H t O 

If heated in air or oxygen, sodium burns with a golden-yellow flame 
to form sodium peroxide. 

2Na-| 0,-»-Na 2 O s 
Action of sodium on water 

Place a small piece of sodium (a piece the size of a very small pea 
will be ample) on the surface of water in a large dish or trough. The 
sodium will dart about and melt to a silvery ball of molten sodium, 
liberating hydrogen and forming sodium hydroxide (see p. 196). 

2Na + 2H g O -»- 2NaOH + H, 




If a light is applied the hydrogen given off will burn with a goMen- 
yellow flame. Sodium and potassium are so readily attacked by the 
oxygen and water-vapour of the atmosphere that they are usually 
kept below the surface of petroleum oil. 

Flame coloration 

Sodium compounds impart a persistent golden-yellow coloration 
to the flame (see flame-test, p. 374). This colour is invisible when 
viewed through blue glass. The above serves as a very definite and 
delicate test for the presence of sodium in a compound. 

Uses of sodium 

Sodium is used in the manufacture of sodium cyanide, NaCN, 
which is employed in the extraction of gold, and also in manu- 
facturing sodamide, NaNH., and sodium peroxide, Na s O,. An alloy 
of sodium and lead has been employed in the manufacture of lead 
tetraethyl, an anti-knock additive used in petrol. 

Pb + 4Na + 4C,H 6 C1 ->• 4NaCl + Pb(C,H 8 ) 4 
The alloy is heated in ethyl chloride vapour. Sodium vapour lamps 
are well known for the intensely yellow illumination they give. 
Sodium has also found a recent use in the reduction of titanium 
chloride to the metal by heat. Titanium is used in high-temperature 


TiCl 4 + 4Na-»- Ti + 4NaCl 



Calcium occurs abundantly and very widely distributed as 
calcium carbonate, CaCO a , which is found as chalk, limestone, 
marble, calcite, Iceland spar and aragonite. 

It also occurs as calcium sulphate as gypsum, CaS0 4 . 2H,0, and 
as anhydrite, CaS0 4 . 

It occurs, less abundantly, as fluorspar, calcium fluoride, and as 
calcium phosphate. 


Calcium is obtained by the electrolysis of fused calcium chloride. 
(Notice that whereas potassium is obtained from the hydroxide, 
calcium cannot be obtained from the hydroxide by electrolysis. The 
reason is that if calcium hydroxide is heated strongly it becomes 
quicklime and, if that is heated strongly, it merely becomes white hot 
and does not melt. Quicklime is one of the most refractory sub- 
stances known.) 



Calcium chloride, which is obtained as a by-product from the 
Solvay process (see p. 3 10), is placed in a graphite (carbon) container 
which is the anode of the cell (Fig. 153). The cathode is an iron rod 
which just dips below the surface of the calcium chloride. The calcium 
chloride is melted and the electrolysis is begun. As the calcium forms 


+ Anode 


FlO. 153. 
Extraction of calcium. 

on the iron rod, the latter is withdrawn and an irregular stick of cal- 
cium is gradually formed. Chlorine is evolved at the anode. If the 
anode were made of metal it would be attacked by the chlorine, but 
chlorine has no effect on carbon, and the gas is a valuable by- 

At the cathode At the anode 

Ca»+ + 2e~ -»- Ca 2C1" - 2e~ -»- CI, 

(a reduction) (an oxidation) 

Physical properties of calcium 




Silvery, shining metal which rapidly 

tarnishes in air owing to formation of 

film of oxide. 


1.55 g/cm-». 


Malleable and ductile. 

Tensile strength 

Possesses fair tensile strength. 



Conduction of heat 

and electricity 

Good conductor of heat and electricity. 

Action of air on calcium 

Calcium is not as reactive as sodium and potassium, and it is not 
necessary to keep it below the surface of petroleum. 



A white film of oxide is formed on the surface on exposure to air. 
Calcium will burn with a brick-red flame if heated in the air and 
forms quicklime, calcium oxide. 

2Ca + O s -*■ 2CaO 

Action of calcium on water (see p. 196) 

Calcium is attacked by water, liberating hydrogen and forming a 
solution of calcium hydroxide. 

Ca -f 2H s O -> Ca(OH) g + H, 

The action is not so vigorous as that of sodium or potassium and a 
test-tube full of water can safely be placed over a piece of calcium in 
a dish containing water, and the hydrogen can be collected. 

Flame coloration (see p. 374) 

Calcium compounds (especially the chloride) colour the flame 
brick red. 


Calcium is used as a dcoxidiscr for steel castings and in the 
extraction of thorium by heating its tetrachloride with calcium. 

ThCl 4 + 2Ca -> Th + 2CaCl a 



Compounds of this metal are quite abundant. Some of the better 
known are: 

Mica, felspar, K 2 AljSi g O, 9 

Kaolin (china clay), Al 2 Si 2 0,.2H 2 (used in making porcelain) 

Corundum, A1 2 3 

Cryolite, Na 3 AIF, 

Bauxite, AU0 3 .2H 8 


Purified bauxite (alumina, Al a 3 ) is electrolysed (Fig. 154) in 
solution in molten cryolite. If more purified bauxite is added as 
required, cryolite is unchanged and can be used indefinitely. Pure 
aluminium is tapped from the cathode, but the carbon anode tends 
to be oxidised away by oxygen. 

2A1 3+ + 6e- -»- 2AI ; 30 2 ~ - 6e~ -»- 1 iO a 
at at 

cathode anode 

Cryolite is sodium aluminium fluoride. 



Crust of solid 

+, carbon anode 

pure aluminium oxide 
in molten cryolite 

carbon lining 
as cathode 

steel trough 

Molten aluminium 

Fig. 154. 
Extraction of aluminium. 

Physical properties of aluminium 





Tensile strength 



Silvery white. 

2.69 g/cm- 3 . 

Can be rolled into foil. 

Moderate (high in alloys). 


Conduction of heat and electricity Good. 

Action of aluminium with air 

The metal acquires a continuous, very thin coating of oxide and 
this resists further action. At 800°C, it will burn in air, forming its 
oxide and nitride. 

4A1 + 30 2 
2AI + N,- 



Action of aluminium with acids 

The metal attacks dilute hydrochloric acid slowly and the con- 
centrated acid rapidly, liberating hydrogen. 

2A1 + 6HCI -*- 2A1C1 3 + 3H, 
Aluminium has no action with dilute sulphuric acid, but the hot, 
concentrated acid is attacked by it with liberation of sulphur dioxide. 

2A1 + 6H 8 SO, -»■ A1 2 (S0 4 ) 8 + 6H.0 + 3S0 8 
Nitric acid does not react with aluminium at any concentration. This 
is probably because it produces on the metal a thin layer of insoluble 
oxide, which protects the metal from further attack. (See passive 
iron, p. 488.) 



Action of aluminium with caustic alkali solution 

The metal, especially in powder form, reacts violently with bench 
sodium hydroxide solution, liberating hydrogen and leaving sodium 
aluminate in solution. (Potassium hydroxide similar.) 

2A1 + 20H- + 6H 8 -► 2Al(OH)," + 3H 8 

Uses of aluminium 

(1) In alloys. The metal is a constituent of several light alloys. 
They combine high tensile strength with lightness, and have been 
much used in aircraft construction. Duralumin (Al, Mg, Cu, Mn), 
magnalium (Mg, Al) and aluminium bronze (Al, Cu) are well known. 

(2) In cooking utensils. Cheapness, low density, good appearance, 
good conductivity for heat and resistance to attack by cooking 
solutions have combined to make aluminium very popular in the 
kitchen. Aluminium vessels must not be exposed to alkaline solutions 
(see above). 

(3) In overhead electric cables. The lightness of aluminium is very 
favourable here. Thick cables of low resistance can be employed 
without undue weight. 

(4) In aluminium paint. The powdered metal is used, with oils. 

(5) In thermit processes. The reactions between aluminium powder 
and oxides of other metals are commonly very exothermic. If a 
mixture of iron(IH) oxide and aluminium powder, known as 'ther- 
mit', is 'fired' by burning a piece of magnesium ribbon stuck into it, a 
violent reaction will occur. Molten iron is produced with a slag of 
aluminium oxide floating on it. 

Fe 2 8 + 2A1 -»- 2Fe + A1 2 8 
This reaction has been used in welding steel parts in situ by means 
of the molten metal produced, and in incendiary bombs in the early 
part of World War II. Similar reactions are also used in isolating 
certain metals, e.g., chromium. 

Cr a 3 + 2A1 -> 2Cr + A1 2 3 


The four metals already described are obtained from their com- 
pounds by electrolysis. We now come to those metals which have been 
obtained for many years and which are used in far greater quantities 
than the more electropositive elements. 

The metals zinc, iron, lead and copper are found chiefly as the 
impure carbonates and sulphides, and iron is also found as the 
impure oxide. The following processes are amongst those commonly 



used in the extraction of metals from their ores although all the 
processes are not used in connection with each individual metal. 

Concentration of ores 

Very often, as ores are found contaminated with earthy impurities, 
methods are employed to pick out the richer ores, or those worth 
working up, and to reject the poorer grades. This may be done by 
hand, or the earthy matter may be washed away by means of a stream 
of water, leaving the heavier ores. In the case of copper ores, the 
latter drop through a magnetic separator where metallic ores are 
deflected into one pile, whilst the lower-grade ores and earthly im- 
purities are not deflected and pass straight on. 

Roasting in air 

Since many ores contain either the sulphide or carbonate of the 
metal, a preliminary roasting in air will remove the sulphur as sulphur 
dioxide and drive off carbon dioxide from the carbonate. Thus: 

2ZnS(c) + 30 2 (g) -* 2ZnO(c) + 2SO,(g) 
ZnCO a (c) -> ZnO(c) + CO z (g) 

The oxides are usually easier to deal with than the sulphides or 

Reduction process 

The roasted ore must now be reduced. The reduction in the case of 
zinc and iron is by means of carbon. In some cases (e.g., copper), by 
a suitable adjustment of the roasting process, it is possible to oxidise 
some of the sulphide to oxide, and then by adding more of the ore to 
supply sufficient sulphide to react with the oxygen of the oxide, 
leaving the metal. 


The product of the reduction process is seldom a pure specimen of 
the metal. Purification may be carried out electrolytically (as in the 
case of copper and zinc). By electrolysis a very pure product is 
usually obtainable. In other cases the impure metal is heated in a 
hearth open to the air, when impurities oxidise and rise to the surface 
as a scum and can be removed. 



Zinc occurs in various parts of the world, as 
Zinc carbonate, ZnCO a , calamine, and 
Zinc sulphide, ZnS, zinc blende. 



The ores are first roasted in air when the oxide is formed whether 
the ore is calamine or zinc blende. 

ZnCOs-^ZnO + CO, 
2ZnS 4- 30 3 -*■ 2ZnO + 2SO, 
The sulphur dioxide is frequently used for the manufacture of 
sulphuric acid. 

The ore is now mixed with coke and placed in a fireclay retort to 
the end of which there is attached a fireclay condenser (Fig. 155). On 




Furnace wall 

Fio. 155. 
Extraction of zinc. 

the end of this condenser is placed an iron 'prolong' which collects 
any zinc which escapes the condenser. The mixture is heated by 
means of producer-gas for about twenty-four hours. The zinc oxide 
is reduced to metallic zinc, the carbon becoming carbon monoxide, 
which burns at the mouth of the condenser. 

ZnO + C -► Zn(g) + CO 
zinc carbon zinc carbon 
oxide monoxide 

The zinc distils out of the retort and the bulk of it condenses to 
molten zinc in the condenser, and is removed from time to time. 
Owing to the presence of air in the retort, some of the zinc burns to 
zinc oxide and condenses on the upper part of the condenser as 'zinc 
dust' (this is a mixture of zinc and zinc oxide). The impure zinc 
obtained in this way is purified by electrolysis. It is frequently 
'granulated' by running the molten metal into water. (In this granu- 
lated form the zinc offers a larger area for action with, e.g., dilute 
acids.) See also p. 155. 

Physical properties of zinc 





Tensile strength 


Bluish-white metal. Can be polished. 

7.1 g/cm" 3 . 

Malleable at temperatures between 100° 

and 150°C. 
Possesses high tensile strength. 



Melting-point 419°C. 

Conduction of heat 
and electricity Good. 

Action of zinc on exposure to air 

Zinc is only very slightly attacked by air owing to the formation of 
a film of oxide which prevents further action. If heated in air, it will 
burn with a bluish-green flame forming the oxide. 

2Zn + O a -> 2ZnO 

► ZnS0 4 + H 2 

zinc hydrogen 

Action of acids 

Ordinary samples of zinc are attacked readily by .the mineral acids. 
Very pure zinc is only slowly attacked, and in recent years the purity 
of even the commercial zinc is so high that very often the action of 
dilute sulphuric acid on zinc is slow at first. The common actions are 
expressed by the equations: 

Zn + H s S0 4 - 
zinc dilute 

With hot concentrated sulphuric acid, sulphur dioxide is formed. 

Zn + 2H 2 S0 4 -»- ZnSO, + 2H 2 + SO, 
Zn + 2HC1 ->- ZnCl a + H s 
zinc dilute zinc hydrogen 

hydrochloric chloride 

3Zn + 8HN0 3 -»- 3Zn(N0 3 ) 2 + 4H a O + 2NO 

(In the last case other oxides of nitrogen and even ammonia, which 
combines to form ammonium nitrate, may be formed.) 

Action of alkalis on zinc 

Zinc is attacked by a hot concentrated caustic alkali solution. 
(Zinc oxide is amphoteric, see p. 259.) Hydrogen is evolved and 
sodium or potassium zincate solution is left. 

Zn + 20H- + 2H g O -*- Zn(OH) t *~ + H, 

Action of water on zinc 

Water does not attack zinc to any appreciable extent. Zinc at a red 
heat is attacked by steam with the formation of hydrogen. 

Zn + H 2 -> ZnO + H, 



Uses of zinc 

Galvanising. Small iron objects and iron wire or sheet are often 
coated with zinc (i.e., galvanised) to delay rusting. This may be done 
by spraying, electrolytic deposition or dipping into molten zinc. Zinc, 
in air, acquires a coherent, inert oxide layer by which rusting of the 
iron is prevented. Also, some protection is still given to the iron even 
if the zinc layer is broken. This is so because zinc is more electropositive 
than iron and the first stage of oxidation of zinc, Zn — *■ Zn i+ + 2e~, 
occurs in preference to that of iron, Fe — *■ Fe* + + 2e~. 

Thin iron sheet can also be tin-plated by passing the cleaned sheet 
through molten tin with a flux present, e.g., zinc chloride. Tin-plate 
is used in canning fruit, meat and fish (to which it imparts no taste). 
A continuous layer of tin protects iron from rusting (being unreactive 
with air). Tin, however, is less electropositive than iron, so, as soon 
as the tin layer is broken, oxidation of iron begins. 

Zinc is also used in alloys, e.g., brass (copper and zinc, 2 : 1) and 
in dry Leclanche batteries, as the negative pole. 



The occurrence of iron, a metal of immense importance, has had a 
profound effect upon the development of the countries in which it is 
found. Your geography book will tell you how important it has been 
to this country that iron ore and coal (which is necessary for the 
extraction of large quantities of the metal) have been found com- 
paratively close together. The chief ores are the following: 

Hamatite, found in Great Britain, United States, France, Germany, 
Belgium and Spain, is impure iron(III) oxide, Fe.0 3 . 

Magnetite, or magnetic iron ore, Fe 3 4 , occurs in Sweden and 
North America. 

Spathic iron ore, ironfU) carbonate, FeCO s , is found in Great 

Iron is widely diffused and is present in many soils. 


Since the demand for iron is so great (345 million tonnes of pig iron 
are made yearly), iron has often to be made from poorer grade ores, 
containing a certain amount of earthy impurities. The ores are first 
roasted in air when iron(lll) oxide, Fe t 3 , is the main product. 

This ferric oxide is mixed with coke and limestone and introduced 
into a blast furnace (Fig. 156). The blast furnace is a tall structure 
about 30 m high and 9 m in diameter at the widest part. It contains 
a firebrick lining inside a steel shell, and a blast of hot air can be 





introduced low down in the furnace through several pipes known as 
tuyeres. A well at the bottom of the furnace serves to hold the molten 
iion and slag until these can be run off. The mixture of ore, coke and 
limestone is fed in continuously from the top, and a blast furnace, 

ferric oxide, 

coke and limestone 

waste gases 
to heat up 
incoming air 


tapped here 

tapped here 

Fio. 156. 

The blast furnace. 

once started, is kept going for months at a time until repairs are 
necessary or work lacking. 

Chemistry of the action 

As the hot air comes into contact with the white-hot coke, the 
latter burns to form carbon dioxide. 

C + O, ->■ CO, 
carbon carbon 

(col») dioxide 

The above reaction liberates a very large quantity of heat, and it is 
this heat which keeps up the high temperature necessary for the 
reduction process. 

As the gas is forced higher up the furnace the supply of oxygen 
(from the air) becomes less and the carbon dioxide coming into con- 
tact with white-hot coke is reduced to carbon monoxide. 
CO, + C -> 2CO 

This carbon monoxide at the high temperature (about 1000°) 
reduces the iron(HI) oxide to metallic iron forming carbon dioxide. 

Fe,O s + 3CO -*- 2Fe + 3CO,. 

irorKIIQ carbon iron carbon 

oxide monoxide dioxide 

The molten iron runs to the bottom of the furnace. 

Action of the limestone. The limestone, which has been introduced 
together with the ore, is first decomposed at this high temperature to 
form calcium oxide. 


»- CaO + CO, 

calcium carbon 
oxide dioxide 

The earthy impurities contain a certain amount of silica (SiO,), 
which is an acidic oxide, and this combines with the basic oxide, 
calcium oxide, to form calcium silicate. 

SiO, + CaO -► CaSi0 3 
silica calcium silicate 

(compare CO, + CaO -> CaCO s ) 

The earthy impurities and this calcium silicate form a molten slag 
which does not mix with the iron but floats above it and can be run 
off separately. At one time this slag was a waste material, and the 
countryside has been defaced by the presence of huge slag heaps. The 
slag is being increasingly used at the present time for making roads. 
The molten iron is run off into moulds, where it is allowed to cool 
in long bars about 1 m long and 10 cm in diameter. It is known as 
'cast iron' or 'pig iron'. 

Cast iron 

This is impure iron and contains varying amounts of impurities, 
such as carbon (4%), with smaller quaiitities of silicon, phosphorus 
and sulphur. This impure iron melts at a lower temperature than pure 
iron and is brittle. It cannot be welded, and possesses little tensile 
strength. (Thus it could not be used for bridges or motor-car con- 
struction.) It is, however, used extensively for small castings, such as 
fire-grates, railings, hot-water pipes, Bunsen burner bases and for 
many other purposes where little strain is imposed. 



Wrought iron 

This is the purest form of iron, and is obtained from cast iron by 
heating it with iron(III) oxide in a furnace by a process known as 
'puddling'. The oxygen of the iron oxide oxidises the impurities, 
carbon and sulphur, to the gaseous oxides which escape, and phos- 
phorus to phosphates and silicon to silicates, and these form as a 
slag. The semi-molten mass is then hammered and rolled so that the 
slag is squeezed out and a mass of almost pure iron remains. 

Wrought iron has a higher melting-point than cast iron. It is 
malleable and can be forged, hammered and welded when hot. It is 
tough and fibrous, and can withstand some strain, but is not elastic, 
and, if subjected to great strain, it will bend. It cannot be tempered! 
It is used to make iron nails, sheeting, ornamental work, horse-shoes, 
and agricultural implements. It has been replaced to a large extent 
in recent years by mild steel, which can be made more cheaply. 


Ordinary steel is a material containing iron and a small proportion 
of carbon, the proportion being determined by the intended use of 
the steel. About 90% of the pig iron made is converted into steel. 

Bessemer process. The usual Gilchrist-Thomas version of the 
Bessemer process is as follows. The Bessemer converter is made of 
steel plate lined with firebrick. It is roughly egg-shaped with an open 
top and a base perforated to take an air-blast. It has a basic lining 
of calcined dolomite (CaO and MgO). White-hot cast iron is run in 
and an air-blast is applied. The various impurities of the cast iron 
oxidise— carbon to its gaseous oxides (CO and COj) which escape, 
manganese and silicon to oxides which form a slag, and phosphorus 
to its pentoxide, P 4 O 10 , which is absorbed as a phosphate by the basic 
lining. When the appearance of the flame indicates the end of these 
changes (about thirty minutes), the required content of carbon is 
added to the metal as anthracite. A short air-blast is applied for 
mixing and the steel can then be poured. 

Recently, the Linz-Donawitz (L-D) modification has been intro- 
duced. The converter is similar but with a solid base. Instead of an 
air-blast from the base, an oxygen blast at 10 atm pressure is applied 
over the top of the white-hot cast iron, by a water-cooled copper tube. 
The rest of the process is essentially the same as before. Lime may be 
added to assist formation of slag. Advantages of the L-D process are 
greater speed and a less brittle steel, because the oxygen blow con- 
tains no nitrogen. Steel from this process is very suitable for pressing 
motor-car bodies. 

Siemens-Martin open hearth process. The chief impurities of cast 
iron are non-metals, the oxides of which are acidic. The object of the 

IRON 487 

Siemens process is to oxidise the impurities to their acidic oxides, and 
these combine with a basic lining supplied to the furnace. In this 
process pig iron, scrap iron and iron oxide (the last of these supplies 
the oxygen for oxidation of impurities) are melted in an open hearth 
which has been lined with a basic material (carbonates of calcium 
and magnesium are used, the oxides being formed at this high tem- 
perature). The amount of carbon is regulated at from 0.5% to 1%, 
and small quantities of various metals such as manganese, nickel, 
chromium or tungsten, are added according to the quality of the steel 
and the use to which the steel will be put 

Properties of steel. Steel is hard, tough and strong. If cooled 
gradually, steel can subsequently be hammered into shape or drilled, 
because it is fairly soft By heating it and suddenly cooling it, the 
steel becomes very hard indeed, of very high tensile strength, and 
elastic. By reheating the steel to carefully regulated temperatures, 
steels of different degrees of hardness and brittleness can be obtained. 
This is called 'tempering'. 

Uses of steel. Ordinary carbon steel, and alloy steels, have extensive 
and well-known uses, e.g., as armour plate in warships and military 
tanks, as pressed sheet in automobile bodies, as girders and wire 
mesh in ferroconcrete building, as cutting and boring tools, crushing 
machinery and stainless cutlery. The following are a few examples: 

with chromium (1-4%) 
with chromium (10-12%) 
with cobalt (2-4%) 
with tungsten (5-18%) 

Physical properties of pure iron 

armour plating, gears 
stainless steel 
in electromagnets 
drills and cutting tools. 



Pure iron is a white metal and can be 

7.9 g/cm- 3 . 
Extremely malleable. 

Tensile strength 
Conduction of heat 
and electricity Good. 

Iron can also be magnetised. 

Action of iron on exposure to air 

In the presence of air and moisture, iron readily rusts, forming a 
reddish-brown solid which consists mainly of hydrated iron(IIl) oxide, 



(Fe a 3 .xH a O). If finely divided (for example iron filings), it will burn 
in air or oxygen to form the magnetic oxide of iron, Fe 3 4 . 

3Fe + 20 2 -»-Fe 3 4 

Action of steam on heated iron 

Iron, at a red heat, is attacked by excess of steam, forming 
magnetic oxide of iron (tri-iron tetroxide) and hydrogen. 

3Fe + 4H a O -»• Fe 3 4 + 4H, 

The above action is reversible (see p. 158). 

[Note that if air and water act together on iron, iron(III) oxide is formed, but 
if cither of these substances acts separately the product is tri-iron tetroxide.] 

Action of acids on iron 

1. Dilute sulphuric and hydrochloric acids. Iron is attacked by these 
dilute acids in accordance with the following equations: 

Fe + H 2 S0 4 -»- FeS0 4 + H 2 

Fe + 2HC1 -»• FeCI 2 + H 2 

The iron(II) salt is obtained because the hydrogen which is given 
off during the action is a reducing agent. 

2. Nitric acid. Dilute nitric acid gives a series of complex reactions 
in which oxides of nitrogen and even ammonia are formed. 

Passive state. If a piece of clean iron is dipped into concentrated 
nitric acid there is apparently no action but the iron no longer 
behaves as a piece of ordinary iron; for example, it will not displace 
copper from copper sulphate solution nor is it attacked by dilute 
nitric acid, which normally does attack it. If, however, the piece of 
iron is scratched while in contact with, say, dilute nitric acid, a 
vigorous reaction occurs. This 'passive state' is supposed to be due 
to a protective layer of oxide (Fe 3 4 ) formed on the iron by the 
strong oxidising agent, concentrated nitric acid. 

Iron will readily combine with sulphur and chlorine when heated 
with them to form iron(II) sulphide (p. 7) and iron(III) chloride 
(p. 362) respectively. 

Compounds of iron(I I ) and iron (ID) 

An iron atom possesses 26 electrons. The shell grouping normally 
shown for them is 2, 8, 14, 2. The 2 adjacent to the 14 are the valency 
electrons and the atom can part with them in electrovalent com- 



bination; it is then said to act in the iron(II) (formerly called ferrous) 
state and produces the ion, Fe a+ . 

If, however, the iron(II) ion is exposed to suitable oxidising con- 
ditions, it will utilise a further electron for valency purposes. The 
resulting ion is then said to be in the iron(III) (formerly called 
ferric) state as Fe 3 " 1 ". 

Fe 2+ -*■ Fe 3 * + er 
The following table shows the formula (in molecular form) of the 
more important simple compounds of iron. 



Valency of iron = 2 

Ion Fe** 




FeS0 4 
Soluble iron(II) com- 
pounds give green sol- 


Valency of iron = 3 
Ion Fe>* 

Fe 2 3 

Fe(OH) 3 

FeCl 3 

Soluble iron(III) com- 
pounds give yellow or 
brown solutions. 

Solutions of pure iron(II) compounds are distinguished from those 
of pure iron(III) compounds by the colour difference just mentioned, 
though, in dilute solution, the green colour of the iron(III) salts is 
very pale. Iron(III) hydroxide and iron(III) oxide are both brown, 
while iron(ll) hydroxide, as usually precipitated, is green. IronfJI) 
oxide is so readily oxidised by oxygen of the air that it cannot be kept 
under ordinary laboratory conditions. 

A simple test for an iron{II) salt. Dissolve a little iron(II) sulphate 
in water. To the solution, add sodium hydroxide (caustic soda) solu- 
tion. A dirty-green, gelatinous precipitate of iron(II) hydroxide is 
formed. This reaction is typical of a ferrous salt. 

Fe* + (aq) + 20H"(aq) -> Fe(OH) 2 (c) 
Where it is exposed to the air, the precipitate will become brown 
because it is oxidised to iron(III) hydroxide. 

2Fe(OH) t + *0 2 + H g O ->- 2Fe(OH) 3 

A simple test for an iron(Ill) salt. Using iron(III) chloride solution, 
repeat the test just given. In this case, the precipitate is reddish brown 
and is iron(III) hydroxide. This reaction is typical of an ironfJII) salt. 
Fe 3, (aq) + 30H"(aq) -*■ Fe(OH) 3 (c) 

Conversion of an iron(ll) salt to an iron(lll) salt. The conversion of 



an iron(Il) salt to an iron(III) salt is an oxidation and is brought about 
by oxidising agents. 

Fe»+ — e- —*■ Fe 3+ 

To a solution of iron(II) sulphate, which is green, add dilute 
sulphuric acid. Warm the mixture and add cautiously a few drops of 
concentrated nitric acid. (A dark brown coloration will probably 
appear. For an explanation of this see the 'brown ring' test, p. 436.) 
Heat the mixture. Brown fumes of nitrogen dioxide are given off and 
a brown or yellow solution remains. It contains iron(III) sulphate 
(test as described above). 

The nitric acid has oxidised the iron(II) sulphate to iron(III) 
sulphate and has itself been reduced to nitrogen monoxide, which, 
on exposure to air, gives nitrogen dioxide. 

6FeS0 4 + 2HN0 3 + 3H 2 S0 4 -+• SFe^SO,), + 4H 2 + 2NO 
or 6Fe* + + 8H + + 2NO s - -»- 6Fe 3+ + 4H 2 + 2NO 

Iron(II) chloride is converted by the oxidising agent, chlorine, to 
iron(lll) chloride, in solution or when heated. 


2FeCI 3 + Cl 2 
2Fe 2+ + CI, 

2FeCl 3 
2Fe s+ + 2CI- 

Conversion of an iron(IU) salt to an iron(II) salt. The conversion of 
an iron(III) salt to an iron(II) salt is a reduction, and is brought about 
by reducing agents. 

Fe» + + e--* Fe* + 

To a solution of yellow iron(III) chloride, add hydrochloric acid 
and zinc. There is vigorous effervescence with evolution of hydrogen. 
Leave the mixture for 20 to 30 minutes. The colour of the liquid is 
now green. It contains iron(II) chloride. (Test, after filtering, as 
described above.) 

The reduction here used to be ascribed to 'nascent' hydrogen, pro- 
duced by the action of zinc and acid, used at the instant of its pro- 
duction and thought to be more active than ordinary gaseous 
hydrogen, which does not reduce an iron(III) salt in solution. 

2HC1 + Zn -> ZnCl, + 2H (nascent) 
2FeCl 3 + 2H -*■ 2FeCl, + 2HC1 

The modern (and better) explanation is that, in the presence of the 
acid, zinc supplies electrons which are taken up by the iron(III) ions, 
which are thereby reduced to iron(II) ions. 

Other reducing agents will convert iron(III) salts to iron(II) salts. 

LBAD 491 

The action of two common ones is represented in the following 

/2FeCl 3 + H 2 S -*- 2FeCI, + 2HC1 + S 
i or 2Fe»+ + S«- -> 2Fe* + + S 
Fe 2 (SO«) a + S0 2 + 2H 2 -»■ 2FeS0 4 + 2H,S0 4 
i or 2Fe 3+ + SO, + 2H s O -> 2Fe» + + S0 4 »- + 4H+ 



Lead occurs as galena, lead sulphide, PbS, and is distributed widely 
in the earth's crust, being found to some extent in most parts of the 
world. It has been known for a very long time— lead pipes were used 
by the Romans in this country. 


The galena is roasted with excess of air to form lead(II) oxide. 

2PbS + 30 a -»-2PbO + 2S0 2 

The oxide is then reduced to lead by heating with carbon in a 

small blast furnace. 

PbO + C->Pb + CO 

Some iron is added to reduce any remaining galena, 
PbS + Fe-»-Pb + FeS 
and lime to combine with earthy impurities and form a molten slag. 
The molten iron(II) sulphide and slag are tapped off separately from 
the lead. 

New process for joint extraction of zinc and lead 

Substantial supplies of ores are known in which lead and zinc 
sulphides occur together. A new plant (in production in England in 
1968) extracts both metals in a single process. 

The ores (previously concentrated by flotation) are roasted in 
excess of air to form oxides. 

2ZnS + 30 2 -*■ 2ZnO + 2SO, (PbS similar) 
The oxides are dropped, with coke (reducing agent), into a blast 
furnace (p. 484) and are reduced to the corresponding metals. 

ZnO + C-*Zn + CO; PbO + C -»• Pb + CO 
Molten lead settles to the base of the furnace, with some slag, and is 
tapped off periodically. Zinc is vaporised and leaves the furnace by 
a pipe near the top, together with nitrogen and oxides of carbon. 
Zinc vapour is stripped from the gases by a spray of molten lead. The 



mixture of zinc and lead flows into shallow tanks where molten 
zinc separates above the much denser molten lead and can be tapped 
off separately. The lead is returned to the stripping circuit The 
process is continuous and economical in both manpower and heating 
costs; a disadvantage is that 400 tonnes of molten lead must be 
pumped round per tonne of zinc extracted. 

Isotopes of lead 

Lead is a product of radioactive change. As such, it is found in 
association with natural deposits of uranium and thorium, both of 
which are radioactive. Uranium (U = 238) is converted to lead by 
the loss of eight alpha-particles, each of mass 4, from its nucleus. 
This gives a lead isotope, Pb = 206. Thorium (Th = 232) loses six 
alpha-particles per atom and gives a lead isotope, Pb = 208. Ex- 
perimental tests have shown the correctness of these two atomic 
weights for samples of lead extracted from uranium and thorium 
minerals. The heavier isotope has two additional neutrons on the 
atomic nucleus, but both isotopes have the same number of nuclear 
protons (82) and the same electron arrangement, so they are chem- 
ically identical. Ordinary lead is an isotopic mixture, Pb = 207.2. 

Physical properties of lead 


Tensile strength 
Conduction of heat 
and electricity. 


Bluish white. 


Very malleable. Can be cut with a knife. 

Has a metallic lustre but speedily 


Good conductor of heat and electricity. 

Action of air and water on lead 

Lead is attacked by air and water together, a white layer being 
formed on the lead which consists of a mixture of lead hydroxide and 
lead carbonate. Lead is used extensively as piping to carry water 
supplies and cases of poisoning have been traced to the removal of 
this layer of hydroxide and carbonate by the water passing through. 
If the water is slightly 'hard' (see p. 201) a protective coat appears to 
be formed and none of the lead is removed. In many water supplies 
nowadays, the water is specially hardened by the addition of lime to 
prevent any of these poisonous effects. 



If lead is strongly heated in air it forms litharge (a yellow powder) 
which is a form of lead(II) oxide 

2Pb + OJ-+ 2PbO 

but if heated to a carefully regulated temperature of about 450°C red 
lead is formed. 

3Pb + 20 a -► Pb 3 4 
(red lead) 

Action of acids on lead 

Dilute sulphuric acid and dilute hydrochloric acid have no action 
on lead. 

Hot concentrated sulphuric acid attacks lead (compare copper): 

Pb + 2H t S0 4 -► PbS0 4 + 2H a O + SO a 
lead(II) sulphur 

sulphate dioxide 

Nitric acid attacks it forming nitrogen monoxide (chiefly). 

3Pb + 8HNO s -► 3Pb(N0 3 ) 2 + 4H s O + 2NO 
dilute nitric lead(II) water nitrogen 

acid nitrate monoxide 

The only satisfactory laboratory method of obtaining 
lead in solution is by the action of dilute nitric acid to form 
lead nitrate solution. 

Uses of lead 

(1) Lead is used extensively in the manufacture of water and gas 
piping and as lead sheet for roofing. It is particularly valuable for 
piping as it is easily repaired and joints are quickly made. It is also 
soft and bends easily at corners. It is also used in the manufacture of 
lead shot, solder, pewter and type-metal. 

(2) For paint. If lead is exposed to the action of acetic acid, air, 
water-vapour and carbon dioxide, it is converted into a basic car- 
bonate known as white lead (Pb(OH) s .2PbCO a ) and as such is used 
for paint. 

(3) Lead is used in the manufacture of electrical accumulators and 
as a covering material for cables. 



Native copper or boulder copper is found as the element round the 
shores of Lake Superior but most of the copper used in industry 
occurs as copper pyrites, CuFeS a . 





Extraction from boulder copper 

This is merely refined, very often on the spot, by building a con- 
tainer round the copper, filling the container with copper(II) 
sulphate solution and making the impure native copper the anode of 
a cell (Fig. 157). A strip of copper is made the cathode and pure 


Pure Co, 

ire Copper 
deposit here 

Boulder Copper 

.Copper Sulphate 

Fio. 157. 
Purification of boulder copper where (he metal is found. 

copper deposits on this from the anode as the electrolysis proceeds. 
(See electrolysis of copper sulphate solution, p. 149). 
At cathode At anode 

Cu a+ + 2e~ — > Cuj. Cu — 2e~ -»- Cu*+ 

Extraction of copper from copper pyrites 

The extraction from sulphur-containing ores is long and tedious 
since the copper must be of a high degree of purity, otherwise it is 
useless for electrical purposes. The chemistry of the process is too 
difficult for inclusion here. 

Physical properties of copper 




A red metal possessing a lustre. It can be 



8.95 g/cm- 8 . 


Very malleable and ductile. 

Tensile strength 

The metal is of fairly high tensile strength. 



Conduction of heat 

It is an excellent conductor of both heat 

and electricity 

and electricity. 

Action of air and water on copper 

Copper is not attacked by pure air or water, but, when exposed to 
the atmosphere, it is slowly attacked on the surface with the forma- 
tion of a green solid. This is probably a basic sulphate. 

When heated in the air, copper forms a layer of black copperfU) 
oxide on the surface. 

2Cu + 2 -»• 2CuO 

Action of acids on copper 

Copper has no action on either dilute sulphuric acid or dilute 
hydrochloric acid. 

With dilute nitric acid, oxides of nitrogen are liberated, chiefly 
nitrogen monoxide, and a blue or bluish-green solution of copper(II) 
nitrate remains. 

3Cu + 8HN0 3 -»- 3Cu(N0 3 ) 2 + 4H 2 + 2NO 
dilute nitric copper(II) water nitrogen 

acid nitrate monoxide 

With hot concentrated sulphuric acid sulphur dioxide is liberated 

and copper(II) sulphate is formed. 

Cu + 2H 2 S0 4 -»- CuS0 4 + 2H 2 + SO a 

Flame coloration 
Copper salts colour the flame a characteristic bluish green. 

Uses of copper 

1. For conducting electric current. It must be very pure because im- 
purities increase electrical resistance. 

2. For ornamental work, being little attacked by the air. 

3. In alloys, e.g., brass (Cu and Zn), bronze (Cu and Sn), German silver 
(Cu, Zn and Ni) and the copper coinage (Cu and Sn). 


Magnesium is extracted by electrolysis of fused magnesium chloride. 
This salt is now prepared from sea-water which contains about 1 million 
tonnes of Mg ,+ per km*. Sodium chloride (see diagram) reduces the 
melting-point of magnesium chloride and allows a lower working tempera- 
ture. Molten magnesium collects and is protected from oxidation by a coat- 
ing of electrolyte and from anodic chlorine by the porcelain sheath. 

At cathode Mg»+ + 2e~ — > Mg; at anode 2C1" - 2e~ — > CI,. 

Properties of magnesium 

Magnesium bums readily when heated in air to produce magnesium 
oxide and nitride. 

2Mg + O, — *■ 2MgO ; 3Mg + N, -»• Mg 3 Nj 
It reacts slowly with boiling water to liberate hydrogen 
Mg + 2H.O -* Mg(OH), + H, 





and burns rapidly when heated in steam (p. 197). All the dilute mineral 
acids react with magnesium at room temperature. Hydrochloric and sul- 
phuric acids liberate hydrogen. 

Mg + 2H+— >-Mg*+ + H, 
When very dilute, nitric acid may liberate some hydrogen but, with ordinary 
dilute acid, a complex reaction occurs with oxides of nitrogen liberated. 
These reactions show the strongly electropositive nature of magnesium. 

Cordon anode 

Iron cathode 
_ (and 
— container) 



fused riaqnesium 
Chloride (with 
Sodium Chloride) 

Fio. 158. 
Extraction of magnesium. 

Uses of magnesium 

1. In several light alloys. For these, see aluminium, p. 479. 

2. As powder, with an oxidising agent such as potassium chlorate, in 
flares and fireworks. 


Electroplating is the electrical precipitation of one metal on another to 
secure improved appearance or greater resistance to corrosion. 

In silver plating, articles such as table-ware or cake dishes, made of 
base alloy, e.g., cupronickcl, are made the cathode in a plating bath of 
potassium (or sodium) argentocyanide, KAg(CN)„ solution. This contains 
some silver ion, Ag+. 

KAg(CN), ^ K+ + Ag+ + 2CN- 
The anode is pure silver. When direct current passes, the following occurs. 
At cathode Ag + + e~ — ► Ag At anode Ag — e~ — *■ Ag+ 
Silver deposits Silver dissolves 

In correct conditions, the silver layer deposited on the cathode article is 
coherent and tough and can be highly polished. 

Chromium plating is much used to improve the appearance of steel 
parts and protect them from rusting. The steel is usually plated first with 
nickel or copper, because chromium does not adhere well on to a steel 
surface. The object is made the cathode in a plating bath which contains 
chromium compounds (e.g., sulphate and oxide) in sulphuric acid and 
water. A lead anode is usual. When direct current passes, chromium 
deposits on the article at the cathode as a bright coherent layer. 

Cr*+ + 3e--*-Cr 
This layer resists rusting and gives a bright 'silvery' appearance. 


1. What are the chief chemical properties which distinguish the metals 
from the non-metallic elements? 

Give three distinct methods by which metals can be converted into their 
oxides, using iron, copper and zinc as examples, and state how the oxides 
of these metals may be distinguished from one another. (B.) 

2 Illustrate, by reference to two chemical and to four physical proper- 
ties the chief differences between the metals and the non-metals. Describe 
briefly how a metal may be isolated from (a) a naturally occurring sulphide, 
and (b) a naturally occurring carbonate. (N.U.J.B.) 

3 Give the chemistry of the extraction of iron from hsmatite (ironflXf) 
oxide). How would you prepare (a) iron(llJ) oxide starting from iron(lll) 
chloride; (b) anhydrous iron(II) chloride starting from iron? (N.U.J.B.) 

4. Name the products, if any, formed when 

(a) air is passed over heated (i) iron; (ii) copper; (in) carbon. 

(b) steam is passed over heated (f) iron; (ii) copper; (ui) carbon. 

(c) hydrogen is passed over heated (i) black iron oxide; (ii) copper 

Give the equations for 
(a") one of these reactions which shows simultaneous oxidation and 

(«■) two of these reactions which together illustrate one reversible 

change. (N.U.J.B.) 

5. Explain what is meant by the statement that iron forms two series of 
salts. Describe how you would prepare from metallic iron one member of 
each of these series, and give one physical and one chemical distinguishing 
test for (a) iron(ll); (/>) iron(III) salts. (B.) 

6. Outline the chemistry of the extraction of (a) zinc from zinc carbonate 
(calamine); (a) lead from lead sulphide (galena). 

State two general methods of preparing salts, and illustrate each method 
by reference to the preparation of a zinc salt. (N.U.J.B.) 

7. The chief ore of zinc is zinc blende, ZnS. Outline the chemistry of the 
extraction of zinc from this ore. State two important uses of zinc. Starting 
from zinc, how would you prepare specimens of (a) zinc oxide and (b) 
crystalline zinc sulphate? (N.U.J.B.) 

8. Describe the electrolytic process for the refining of copper. Name four 
distinct uses of copper. . 

Starting with cupric oxide, how would you obtain (a) a dry crystalline 
specimen of copper(ID sulphate; (6) a sample of metallic copper? 

9. Select, from the following, the decisive reason for regarding (a) zinc 
as a metal, (6) bromine as a non-metal : (a) (i) zinc is much denser than 
water, (ii) zinc conducts electricity well, (iii) zinc is an electropositive 
element, (iv) zinc is malleable at certain temperatures; (b) (i) bromine is a 
liquid at s.t.p., (ii) bromine is an electronegative element, (iii) bromine is 
much less dense than mercury, (iv) bromine combines fairly readily with 

Chapter 37 

Radioactivity; a Brief Account 

RADIOACTIVITY was first noticed and investigated in 1896 by 
Becquerel, taking place in salts of uranium. Becquerel found that 
these salts emitted rays with the following properties. They affected 
a photographic plate in the same way as rays of light, discharged a 
gold-leaf electroscope and caused phosphorescence in certain 
materials, e.g., zinc sulphide. Becquerel did not follow this investiga- 
tion very far, but his discoveries aroused interest and activity in other 
workers. In particular, Marie and Pierre Curie, working in Paris 
about 1900, detected more intense radioactivity in certain mineral 
specimens, such as Bohemian pitchblende, and were able to isolate 
two radioactive elements, to which they gave the names polonium (in 
honour of Marie Curie's native Poland) and radium. Many other 
radioactive elements are now known. 

Kinds of radiation 

Becquerel found that the radiation emitted by uranium was not 
uniform but could be separated into three different types under the 
influence of an electrostatic field. The three types were named alpha-, 
beta- and gamma-rays. 

Alpha-rays were deflected towards the negative plate in the electro- 
static field and so, since unlike charges attact each other, must carry 
a positive charge themselves. The deflection was not very marked, 
from which it was concluded that the moving particles constituting 
the rays were relatively massive. It has since been shown that alpha- 
rays are, in fact, helium ions in rapid motion. Helium ions are formed 
from helium atoms by loss of two electrons per atom. They then 
carry a double positive charge resulting from their possession of two 
nuclear protons and have a mass of about 4 units on the scale of 
"C = 12. 

He -»• He* + + 2e~ 

Beta-rays showed a very marked deflection towards the positive 
plate of the electrostatic field. From this, it was concluded that the 


radioactivity; a brief account 499 

particles of which beta-rays are composed carry a negative charge and 
are relatively small in mass. Beta-rays have since been shown to con- 
sist of electrons in rapid motion. Beta-radiation is simply a moving 
stream of electrons. 

Gamma-rays were not deflected at all by an electrostatic field; that 
is, gamma-rays carry no electrical charge. They are not material 
particles but consist of electromagnetic waves of very short wave- 
length (about 5 X 10 - " cm). That is, they are of the same nature as 
light but of higher frequency. 

Particle of screens 
uranium salt 1/ ^M 


Electrostatic field 


iamma-r aya 


Fig. 159. 

Composite radiation from uranium salt. 

Nature o